Chemistry of Group 14 Elements PDF

1 CHEMISTRY OF GROUP 14 ELEMENTS The elements in Group 14 or IVA of the Periodic Table includes carbon (C), silicon (Si), germanium (Ge), tin (Sn) and lead (Pb): show a gradation from C, which is non-metallic, to Pb, which, though its oxides are amphoteric, is mainly me...

1 CHEMISTRY OF GROUP 14 ELEMENTS The elements in Group 14 or IVA of the Periodic Table includes carbon (C), silicon (Si), germanium (Ge), tin (Sn) and lead (Pb): show a gradation from C, which is non-metallic, to Pb, which, though its oxides are amphoteric, is mainly metallic in nature. Carbon, the extremely distributed element in nature, is an essential constituent of life. Carbon- containing compounds form a special branch of chemistry known as organic chemistry. Carbon also forms many important inorganic compounds, especially the organometallic compounds. Silicon has proved to be an important element in the development of modern technologies such as electronics, construction and other industrial usages. The so-called ‘diagonal line’ through the p- block separates metallic from non-metallic elements and passes between Si and Ge, indicating that Si is non-metallic and Ge is metallic. However, this distinction is not definitive. In the solid state, Si and Ge possess a covalent diamond-type structure, but their electrical resistivities are significantly lower than that of diamond, indicating metallic behaviour. Silicon and germanium are classed as semi-metals. Tin and lead are in use since ancient times. Carbon has two electron and silicon have eight electrons in their penultimate shell but other elements of this group have eighteen electrons in their penueltimate shell. From germanium onwards, there are 10d electrons in the elements and Pb has 14f electrons in its 4f subshell. However, the general outer-shell electronic configuration of these elements can be represented as ns2np2 Electronic structure of Group 14 elements Element Electronic configuration Carbon (6C) [He] 2s2 2p2 Silicon (14Si) [Ne] 3s2 3p2 Germanium (32Ge) [Ar] 3d10 4s2 4p2 Tin (50Sn) [Kr] 4d10 5s2 5p2 Lead (52Pb) [Fe] 4f14 5d10 6s2 6p2 GENERAL PROPERTIES OF GROUP 14 ELEMENTS 1. Atomic and Ionic Sizes The atomic and ionic sizes of these elements are smaller as compared to Group 13 elements and increase down the group from C to Pb. This is due to the effect of addition of extra shell in each element on moving down the group which overcomes the effect of increased nuclear charge. 2 2. Ionisation Energy The first IE of Group 14 elements is larger as compared to Group 13 elements due to the effect of increased nuclear charge. However, on moving from top to bottom in the group, the IE goes on decreasing (except for Pb). The decrease from C to Si is larger as compared to the decrease for Si to Ge and from Sn to Pb. This is due to the reason that from C to Si, the increase in atomic size decreases the effective nuclear charge. While in case of Sn, the presence of extra ten d electrons with poor shielding effect results in increase of effective nuclear charge which decreases the effect of increased atomic size. In case of Pb, the extra fourteen f electrons with even more poor shielding effect increases the effective nuclear charge and its IE increases. 3. Metallic and Nonmetallic Character As the ionisation energy decreases from top to bottom, metallic character goes on increasing. Thus, carbon and silicon are nonmetals, germanium is a metalloid (partly nonmetal and party metal), while tin and lead are distinctly metals. This change from nonmetallic to metallic character is reflected in the physical properties (malleability, ductility and conductance) as well as chemical properties (tendency to form M2+ ions and increase in basic character of oxides and hydroxides). 4. Oxidation States The most common oxidation state of these elements is (+IV) and some compounds are formed in the (+II) state. However, on moving down the group, the stability of (+IV) state decreases while that of (+II) state increases. This is due to the presence of inert-pair effect (as discussed in case of Group 13 elements). Thus, Ge (+IV) is stable while Sn (+II) is a strong reductant. Sn (+IV) is also stable while Sn (+II) is a strong reductant. On the other hand, Pb (+II) is stable and Pb (+IV) is an oxidant. 5. Nature of Compounds Due to presence of ns2np2 electrons, the majority of these compounds are tetracovalent. The covalent nature of the compounds is due to high IE of these elements which does not support the formation of stable M4+ ions. Rather M2+ ions are more stable in case of Sn and Pb. Thus, Sn2+ and Pb2+ ions are ionic as in agreement to Fajan’s rules. Carbon also has a tendency to form ion and C4- ion (in case of some carbides). 3 6. Electronegativity Carbon is the most electronegative out of these elements followed by Si. Further decrease is not apparent due to presence of d-electrons in case of Ge and Sn and f-electrons in case of Pb. 7. Melting Point and Boiling Point Carbon shows the highest melting point and boiling point due to stable lattice with strong M-M bond. The values decreases from top to bottom due to comparatively weaker M-M bond with increased atomic size. 8. Formation of Complexes Complex formation tendency of an element is favored by small atomic or ionic size, high charge and presence of vacant orbitals of appropriate energy. Except C, other Group 14 elements have vacant d-orbital and hence these elements can expand their octet and form complex compounds. Thus, C forms only tetravalent compounds while Si and Ge can form octahedral compounds in the coordination number 6, such as [SiF6]2– and [GeF6]2–. Similarly, Sn and Pb can show a coordination number of 8 in some compounds. 9. Catenation Carbon possesses a remarkable property of linkage of identical atoms to form long chains and rings. This property is known as catenation. The tendency of catenation decreases from C to Pb with decrease in bond energy. ANOMALOUS BEHAVIOUR OF CARBON Carbon differs in many aspects, from the rest of the elements of Group 14, due to its small atomic size, high ionisation energy, and high electronegativity and non-availability of vacant d-orbitals. Some of the characteristic properties are the following: 1. Carbon forms only tetravalent compounds with maximum covalence of 4. 2. It forms pπ-pπ multiple bonds in a number of compounds. 3. Carbon has a marked tendency of catenation and thus forms a large number of compounds. CARBON AND SILICON-COMPARISON OF PROPERTIES Carbon and silicon are the two typical elements of Group 14 with many similar properties and many differences in their properties. Similar Properties 1. Carbon is an essential constituent of organic life (animal and plant kingdom) while silicon is an essential constituent of inorganic materials. 4 2. Both these elements have eight electrons in their penultimate shell (ending with ns2np2 configuration) 3. Both elements form similar compounds such as oxides (CO2, SiO2), hydroxides (CH4 SiH4), halides (CCl4, SiCl4) and oxoacids (H2CO3, H2SiO3). 4. Both elements are typical nonmetals. Different Properties 1. Carbon has higher melting and boiling points than Si. Thus, carbon does not melt, while silicon melts at higher temperatures. 2. Carbon is a good conductor of electricity in some of its forms such as graphite, while silicon is a bad conductor of electricity. 3. Carbon has a high tendency to catenate while Si has a lower tendency to show catenation. 4. Carbon can form pπ-pπ bonds with many elements while Si cannot do so due to its larger size. As a result, CO2 is a gas while silica is a hard solid. 5. Hydroxides of carbon are more stable than those of silicon. 6. Carbon tetrachloride is stable and not hydrolysed in water due to absence of d-orbitals in carbon. On the other hand, silicon tetrachlorideis not hydrolysed in water but also forms addition compounds due to presence of d-orbitals in silicon. OCCURRENCE OF GROUP 14 ELEMENTS The two long-established crystalline allotropes of carbon, diamond and graphite, occur naturally, as does amorphous carbon (e.g. in coal). Diamonds occur in igneous rocks (e.g. in the Kimberley volcanic pipes, South Africa). Carbon dioxide constitutes only 0.04% of the Earth’s atmosphere, and, although vital for photosynthesis, CO2 is not a major source of carbon. During the 1990s, it was discovered that molecular allotropes of carbon, the fullerenes, occur naturally in a number of deposits in Australia, New Zealand and North America. Soot contains fullerenes and related carbon species (nanotubes, concentric fullerenes, open hemi-shells). The formation of soot under fuel-rich conditions involves growth of polycyclic aromatic hydrocarbons, which aggregate to form particles. The current development of the chemistry of fullerenes and carbon nanotubes relies on their laboratory synthesis. Elemental Si does not occur naturally, but it constitutes 27.7% of the Earth’s crust (Si is the second most abundant element after O) in the form of sand, quartz, rock crystal, flint, agate and silicate minerals. In contrast, Ge makes up only 1.8 ppm of the Earth’s crust, being present in trace amounts 5 in a range of minerals (e.g. zinc ores) and in coal. The principal tinbearing ore is cassiterite (SnO2). Important ores of lead are galena (PbS), anglesite (PbSO4) and cerussite (PbCO3). ENTRACTION OF GROUP 14 ELEMENTS Sources of natural graphite are supplemented by manufactured material formed by heating powdered coke (high temperature carbonized coal) with silica at 2800 K. 3C + SiO2 → SiC + 2CO↑ SiC → Si↑ + C Diamond films may be grown using a chemical vapour deposition method and hydrothermal processes are currently being investigated. The manufacture of amorphous carbon (carbon black, used in synthetic rubber) involves burning oil in a limited supply of air. Silicon (not of high purity) is extracted from silica, SiO2, by heating with C or CaC2 in an electric furnace. SiO2 + CaC2 → Si + Ca + 2CO SiO2 + 2C → Si + 2CO Impure Ge can be obtained from flue dusts collected during the extraction of zinc from its ores, or by reducing GeO2 with H2 or C. For use in the electronic and semiconductor industries, ultrapure Si and Ge are required, and both can be obtained by zone-melting techniques. Tin is obtained from cassiterite (SnO2) by reduction with C in a furnace, but a similar process cannot be applied to extract Pb from its sulfide ore but it was extracted by crushing and finely powered ore is sieved and concentrated by froth flotation process. The concentrated ore is roasted in a special reverberatory furnace in a limited supply of air As a result, PbS is partly oxidized to PbO and partly to PbSO4. 2PbS + 3O2 → 2PbO + 2SO2 PbS + 2O2 → PbSO4 After this, more galena is added and supply of air is cut off. As a result, PbO and PbSO4 are reduced to metallic lead which is collected from the bottom of the furnace. 2PbO + PbS → 3Pb + SO2 PbSO4 + PbS → 2Pb + 2SO2 6 CHEMISTRY OF CARBON Physical Properties Carbon exists in different allotropic forms. Its crystalline forms are graphite, fullerene and diamond, while, the amorphous forms are coal, coke, charcoal, lampblack, carbon black, gas carbon and petroleum coke. These forms will be discuss in details later. 13.5.3 Chemical Properties (a) All the allotropic forms of carbon burn on heating in presence of air or oxygen to give CO (in limited supply of O2) or CO2 (in excess of O2). 2C + O2 (limited) → 2CO C + O2 (excess) → CO2 (b) Carbon acts as a reducing agent as supported by the following reactions. Action with acids: C + 2H2SO4 → 2SO2 + CO2 + 2H2O C + 4HNO3 → CO2 + 4NO2 + 2H2O Action with metallic oxides: ZnO + C → Zn + CO PbO + C → Pb + CO Action with sulphates: BaSO4 + 4C → BaS + 4CO Action with water: C + H2O → CO + H2 (c) Carbon combines with nonmetals to give the corresponding compounds. C + 2S → CS2 (d) Carbon forms carbides on heating with lime and silica in an electric furnace. Allotropes of Carbon 1. Diamond Diamond is the hardest and densest natural substance known so far with the highest melting point (4200 K). This is because of its three-dimensional tetrahedral structure consisting of sp3 hybridised C-atoms tetrahedrally linked to four other carbon atoms through strong covalent bonds. Structure of Diamond This results in a large three-dimensional macromolecule with C-C bond length equal to 154 pm. This numerous linkage of C atoms results in its high density, hardness and high melting point. 7 (a) Properties of Diamond (i) Diamond is the purest form of carbon without any shape or lustre. It is polished and cut to enhance its beauty. (ii) It is transparent to X-rays and light and has a very high refractive index (2.417). (iii) It is a bad conductor of heat and electricity due to absence of any unpaired electrons in its lattice. (iv) It burns in air at 1175 K to form CO2. It turns to graphite in vacuum at 2075-2275K. (v) It is slowly oxidised to CO2 when heated with a mixture of K2Cr2O7 and conc. H2SO4 at 475 K. (b) Uses of Diamond (i) Diamonds are beautifully cut and polished to enhance their reflecting property. These are used in jewellery. (ii) Diamonds are used as abrasives and cutting tools due to their hardness. (iii) Diamond dies are used to draw thin wires of metals. 2. Graphite Graphite has a two-dimensional sheet like polymeric layered structure with sp2 hybridised C-atoms linked to three other carbon atoms in hexagonal rings. These layers are about 335 pm apart with the C-C bond length equal to 142 pm. This wide separation results in low density and a weaker lattice. Thus, graphite can be easily split along the lines of the plane. These layers can easily slide over each other due to weaker forces and impart softness, slippery texture and lubricating nature to graphite. Structure of graphite (a) Properties of Graphite (i) Graphite is a dark grey crystalline form of carbon with a metallic and slippery touch. (ii) It is also known as black lead or plumbago as it marks paper black just like lead. It is mixed with required quantities of clay and is used in the lead of lead pencils. (iii) It is a good conductor of heat and electricity due to the presence of one unpaired electron with each C atom. (iv) It is quite inert and difficult to ignite. However, it burns on heating in the presence of air to liberate carbon dioxide. (v) It is heated with conc. nitric acid to obtain a yellow mass known as graphitic acid. 8 (b) Uses of Graphite (i) Suspensions of graphite in water (aqua-dag) and in oil (oil-dag) are extensively used as lubricants. (ii) It is used to make high-temperature refractory crucibles. (iii) It is used in lead pencils and for making stone polish. (iv) It is used as electrodes in electric furnaces. 3. Fullerenes Fullerenes form the family of polyhedral carbon allotropes C2n, (n = 30 – 48), discovered in 1980 with spherical shell structures. The first fullerene molecule discovered was given the name Buckminster Fullerene (C60) due to its resemblence with the geodesic domes developed by Buckminster Fuller. These are obtained when a large electric current is passed through graphite rods in presence of a quenching atmosphere of helium gas. The graphite rods evaporate to give a light, fluffy mass called fullerene soot. It is dissolved in common organic solvents and the different components such as pure C60, C70, C76, C78, C80, C82, C84, C96, etc. are obtained by chromatographic separations. The individual solutions are concentrated to obtained crystals containing solvent and other guest molecules in the interstices. These crystals are sublimed under vacuum to give solvent-free crystals. The most common fullerene, C60 has a roughly spherical structure consisting of arrays of 60 carbon atoms arranged in a truncated icosahedron in which 20 hexagons and 12 pentagons of carbon atoms are fused together to give a football-like geometry as shown in Fig 13.3. In this structure, there are both single and double bonds with C-C bond lengths of 1.453 and 1.383 Å respectively. The structure is very strained and thermally less stable as it starts decomposing at around 750°C, also due to its low heat of formation (42.5 kJ/mol). However, it exhibits great kinetic stability. Structure of Buckminster-fullerene (C60) 9 1. Fullerene mainly shows the reactions of double bonds such as typical additions with amines and substitutions with other reagents as shown below: 2. It can also act as a ligand as shown below: 4. Carbon nanotubes Carbon nanotubes were discovered in 1991 and consist of elongated cages, best thought of as rolled graphene sheets. In contrast to the fullerenes, nanotubes consist of networks of fused 6-membered rings. Nanotubes are very flexible and are of importance in materials science. 5. Coal It is believed to be formed by the slow carbonisation of vegetables matter, in limited supply of air, buried underneath the earth and in presence of high temperature and pressure. Coal is commonly used as a fuel in boilers, furnaces, engines, etc. Bituminous coal is the most common variety of coal, which burns with a smoky flame. It is very hard just like stone. Anthracite contains the highest percentage of C and is use as a reducing agent in various metallurgical operators. It is also used to manufacture synthetic petrol and fuel gases such as producer gas, water gas and coal gas. 6. Wood Charcoal It is a dark greyish black, porous and brittle solid obtained by slow pyrolysis of wood in absence of oxygen. It is an impure form of carbon due to presence of ash. It is heavier than water due to presence of large amount of absorbed air in its pores but when red hot, air is expelled out and it sinks. It can absorb many odouriferous gases and colouring matter and hence is used as a decolorising agent. The extent of adsorption increases with decrease of temperature and increase of boiling points of the 10 gases being adsorbed. It is used as a good fuel and also as a reducing agent in metallurgical operations. It is also used as a constituent of gunpowder. 7. Bone Black or Animal Charcoal It is obtained as a black residue by destructive distillation of bones in a retort. It contains about 10% of carbon and the rest constitutes calcium phosphate in the form of a porous framework. It is a stronger adsorbent as compared to wood charcoal and is more extensively used for decolourisation. On burning, it converts to calcium phosphate, known as bone ash, which is used to manufacture phosphorous and phosphoric acid. 8. Sugar Charcoal It is the purest form of carbon obtained by heating sugar strongly and in absence of air or by treating sugar with concentrated H2SO4. It is a black powdery substance, which contains some organic impurities. It is used for the preparation of artificial diamond and in making gas masks. The adsorption capacity of charcoal can be increase by heating it in a current of superheated steam at 1273 K. This charcoal is known as activated charcoal. 9. Lamp Black or Carbon Black If a carbon-rich compound such as petroleum, kerosene oil, turpentine oil, or natural gas is burnt in a limited supply of air; a sticky soot is obtained consisting of minute particles of carbon, called lampblack or carbon black. It contains about 98 to 99% of carbon and is almost a pure form. It is used for the manufacturing of black ink, typewriter ribbon, black paint, shoepolish, varnishes, etc. It is also used as a kajal and for colouring rubber products. 10. Gas Carbon It is a very pure form of carbon and is obtained as a hard deposit on the sides and roofs of iron retort used for the destructive distillation of coal. It is used in the lining of furnaces and for making electrodes as it is a good conductor of electricity. 11 CHEMISTRY OF SILLICON Properties of Silicon Silicon is a hard and brittle solid which melts at 1693 K. It has similar appearance as that of graphite but has comparatively very low electrical conductivity. Crystalline silicon is chemically less active than the amorphous form. It does not react with oxygen even on strong heating while the amorphous form burns brilliantly and vigorously in oxygen to form the oxide. Si + O2 → SiO2 The amorphous form decomposes steam at red heat and liberates hydrogen. Si + 2H2O → SiO2 + 2H2 It dissolves in hot caustic alkaline solutions to liberate hydrogen and silicates are formed. Si + 2NaOH + H2O → Na2SiO3 + 2H2 Sodium silicate is also formed on treatment with sodium carbonate. Si + Na2CO3 → Na2SiO3 + C Silicon reacts with magnesium and carbon to form silicides at the temperature of the electric furnaces. Si + 2Mg → Mg2Si Silicon burns on heating with F2 and Cl2 to give the corresponding tetrahalides. Si + 2F2 → SiF4 Si + 2Cl2 → SiCl4 It also combines with oxygen, sulphur and nitrogen at elevated Uses of Silicon Ultra pure silicon finds an important place in the electronics industry. The insulating nature of the material is converted to semiconductor after doping and then used in micro-miniaturised devices such as computer chips and transistors. It is mainly used to prepare industrially important alloys such as ferrosilicon and silicobronze. Ferrosilicon is used to increase the acid resistance of steel and for the manufacture of the apparatus used for distillation of strong acids. Silicobronze is used to manufacture telegraph and telephone wires. 12 CHEMISTRY OF GERMANIUM Properties Germanium is a silvery white, brittle but hard metal, which melts at 937°C. It has no allotrope like other Group 14 elements. Germanium is more reactive than silicon, but is stable in air and is oxidised only at red heat to be covered by the protective coating of GeO2. It is not affected by water below 200°C, but it decomposes water at high temperature. Ge + 2H2O → GeO2 + 2H2 It is easily attack by molten alkalis to liberate hydrogen Ge + 2KOH + H2O → K2GeO3 + 2H2 It is not attacked by dil. HCl or H2SO4 but is attacked by dil. HNO3 and conc. H2SO4 2Ge + 2HNO3 → 2GeO2 + 2NO + H2 Uses of Germanium It has remarkable electrical properties, and is use in transistors and resistance thermometers. CHEMISTRY OF TIN Properties of Tin Tin exists in three allotropic forms: the crystalline rhombic and tetragonal or white form and the amorphous grey form with their transitions as follows White tin is the heaviest and the most stable form of tin. It is silvery white in colour and is harder than lead but softer than zinc. It is malleable and ductile at 100°C and is drawn into thin sheets known as tin foil. However, it becomes brittle and powdery at 200°C. It exists between 18°C and 170°C and slowly converts to grey form at its transition temperature. The transition takes place rapidly at 223 K and in presence of a little grey tin, it completely converts to the brittle grey form. The brittle form crumbles to a grey powder known as tin disease or tin pest. The metal produces a cracking noise on bending. This is known as tin cry and happens due to rubbing of metal crystals over one another. It melts at 232oC but boils at very high temperature 2270oC. Tin is weakly electronegative and comparatively less reactive than silicon. Sn2+ + 2e- → Sn; E°= -0.136 V 13 Thus, it is not affected by air and water at ordinary temperatures. However, it burns with a bright flame when heated strongly at 1775-1875 K and forms stannic oxide. Sn + O2 → SnO2 Molten tin decomposes steam to liberate hydrogen. Sn reacts slowly with dilute HCl, but rapidly with hot concentation HCl to liberate a rapid stream of hydrogen. Sn + 2HCl → SnCl2 + H2 Sn does not react with dil.H2SO4, but liberates sulphur dioxide with hot conc. H2SO4. Sn + 2H2SO4 → SnSO4 + SO2 + 2H2O It reacts with cold and very dilute nitric acid to give ammonium nitrate, while nitrogen dioxide is liberated with hot and conc. HNO3. 4Sn + 10HNO3 (cold and dilute) → 4Sn(NO3)2 + NH4NO3 + 3H2O It dissolves in hot alkali to liberate hydrogen. It combines with Cl2 and sulphur on heating. Sn + 2Cl2 → SnCl2 Sn + 2S → SnS2 Uses of Tin Since the metal is not affected by the organic acids or the atmosphere, hence tin is extensively used for tinning of copper, brass and the other metallic utensils. The utensils are cleaned by rubbing with sand and then heated. Tin is also used in tin plating of iron and mild steel. In this process, iron and mild steel sheets are cleaned and dipped in dilute acid so as to remove the oxide film. After this, the sheets are dipped in molten tin and passed through hot rollers to obtain a thin coberent film of tin on the surface. Tin foil is used for wrapping cigrattes etc. Tin amalgam is used in mirrors. 14 CHEMISTRY OF LEAD Properties of Lead Lead is a bluish grey metal which marks paper black. It is highly malleable, but not much ductile and has low tensile strength. It possesses a bright metallic lusture which is lost on exposure to moist air due to the formation of basic carbonate layer on the surface. At ordinary temperature, dry air has no action on its surface. However, on heating it forms litharge which is further oxidised to red lead at high temperature. 2Pb + O2 → 2PbO 2PbO + O2 → 2Pb3O4 Lead dissolves in aerated water and converts to lead hydroxide with appreciable solubility. This phenomenon is known as plumbo solvency (solvent action of water on lead). Plumbo solvency increases with presence of nitrates, organic acids and ammonium salts, but decreases with presence of sulphates, carbonates and phosphates in water due to the formation of an insoluble protective coating of lead salts on the surface of the metal. Plumbo solvency results in the accumulation of lead salts in the body and leads to lead poisoning resulting in death of the victim. Dilute and conc.sulphuric acid and dilute hydrochloric acid results in the formation of an insoluble film of lead sulphate and lead chloride respectively on the lead surface and further action is prevented. Pb + 2H2SO4 → PbSO4 + 2H2O + SO2 However hot conc. HCl results in the formation of chloroplumbic acid and hydrogen is liberated. Pb + 2HCl → PbCl2 + H2 PbCl2 + 2HCl → H2PbCl4 Lead liberates nitric oxide with cold & dilute nitric acid, while nitrogen dioxide is liberated with conc. nitric acid. 3Pb + 4HNO3 (cold & dilute) → 3Pb(NO3)2 + 2NO+ 4H2O Pb + 4HNO3 (conc.) →Pb(NO3)2 + 2NO2 + 2H2O It forms lead acetate with acetic acid in presence of oxygen. Pb + 2CH3COOH + ½O2 → (CH3COO)2Pb + H2O It slowly dissolves in caustic alkalis to form plumbites and liberates hydrogen. Pb + 2NaOH → Na2PbO2 + H2 It reacts with chlorine and sulphur on heating and forms lead tetrachloride and lead sulphide respectively. Pb + 2Cl2 → PbCl4 Pb + S → PbS 15 Uses of Lead Lead is used to make lead containers, pipes, bullets and lead accumulators. It is used for lead coating of telegraph and telephone wires. Due to its resistance towards water and acids, and easy workability, it is used in fine arts. It is used in many alloys. COMPOUNDS OF GROUP 14 ELEMENTS 1. Oxides and Oxoacids All the element of the group form oxides and oxoacids Carbon forms carbon monoxide (CO), carbon dioxide (CO2), carbon surboxide (C3O2), carbonic acid (H2CO3), Peroxocarbonates and Peroxodicarbonates Carbon monoxide is obtained by ignition of carbon in a limited supply of oxygen. It is a colourless, odourless and tasteless gas. It is highly toxic in nature due to its ability to combine with haemoglobin to form carboxy haemoglobin. As a result, the oxygen carrying ability of haemoglobin is replenished and the person inhaling CO dies. Carbon dioxide is produced by burning of carbon or its compounds in excess of oxygen. It dissolves fairly in water to give carbonic acid, hence, it also known as carbonic acid gas or carbonic anhydride. Its gas at room temperature. Liquid CO2 is use in fire extinguishers, as it does not support combustion. In dry powder extinguishers, a mixture of NaHCO3 and sand is used which when thrown over the burning fire, releases CO2 and extinguishes the fire. 2NaHCO3 → Na2CO3 + CO2 + H2O Carbon surboxide (C3O2): It is obtained as a foul smelling, colourless gas by dehydrating malonic acid with phosphorus pentoxide. The suboxide is linear and stable at -78oC. It condenses to a liquid boiling at 6oC, which polymerises to a yellow solid at room temperature and to a red-purple solid on heating to a higher temperature. Carbonic acid is a weak dibasic acid and exists only in solution. Thus, it forms two series of salt normal salts or carbonates, and acid salts or bicarbonates. The peroxocarbonates and peroxodicarbonates are the oxocarbon compounds of the hypothetical peroxocarbonate acids, namely peroxocarbonic acid and peroxodicarbonic acid 16 Sodium and barium peroxocarbonates have been prepared by passing carbon dioxide through the aqueous solutions of sodium and barium peroxide solutions respectively Silicon forms two types of oxides, silicon monoxide (SiO) and silicon dioxide (SiO2). Silicon dioxide (silica) exists naturally as sand and is found in three crystalline forms namely quartz, tridymite and cristobalite, exhibiting low temperature and high temperature polymorphism. It can also be prepared by heating silicon in oxygen. Si + O2 → SiO2 Silica is an unreactive acidic oxide and hence does not react with acid, except HF to form silicon tetrafluoride. SiO2 + 4HF → SiF4 + 2H2O It is reduced to silicon carbide (carborundum), when heated with carbon. SiO2 + 3C → SiC + 2CO The general term ‘silica acids’ refers to the hypothetical compounds which have not been isolated. However, a number of silicates with definite compositions have been obtained. The most important silica acids are orthosilicic acid (H4SiO4) and metasilicic acid (H2SiO3). Orthosilic acid, H4SiO4 or Si(OH)4, is obtained as a gelatinous precipitate by the hydrolysis of silicon tetrafluoride. 3SiF4 + 4H2O → 2H2SiF6 + H4SiO4 The precipitates are washed with water and ether followed by drying between folds of filter paper to obtain a white amorphous powder. Metasilicic acid, H2SiO3 is obtained as a gelatinous precipitate by the treatment of concentrated solution of sodium silicate with hydrochloric acid. Na2SiO3 + 2HCl → H2SiO3 + 2NaCl The precipitates are dehydrated with 90% alcohol and dried to obtain a white amorphous powder. Germanium forms germanium monoxide, GeO and germanium dioxide, GeO2. The monoxide is obtained by dehydration of the hydroxide or by heating the metal in presence of the dioxide at about 1000°C. Ge(OH)2 → GeO + H2O 17 Ge + GeO2 → 2GeO It undergoes disproportionation to give the metal and the dioxide on heating to 700oC. The dioxide is obtained by oxidation of the metal at red heat. It is weakly acidic which reacts with basic oxides to give germanates. Tin forms two types of oxides i.e. SnO, stannous oxide and SnO2, stannic oxide. Stannous oxide is obtained as a dark grey or black powder by heating stannous hydroxide or oxalate in presence of carbon dioxide. It is amphoteric and dissolves in acid to form stannous salts, while stannites are formed in alkalis. SnO + 2HCl → SnCl2 + H2O SnO + 2NaOH → Na2SnO2 + H2O The stannites absorb oxygen and change into stannates. 2Na2SnO2 + O2 → 2Na2SnO3 Stannic oxide is obtained as a white powder by the ignition of tin in air or by calcination of metastannic acid. H2SnO3 → SnO2 + H2O It is insoluble in water and is amphoteric. It dissolves in halogen acids to give hexahalostannates, while stannic sulphate is formed with dil. H2SO4. However it is not affected by other acids. Lead forms lead monooxide (PbO), lead dioxide (PbO2), Red lead (Pb3O4), lead suboxide (Pb2O) and lead sesquioxide (Pb2O3). Lead monoxide is obtained by heating lead in air at about 575k. 2Pb + O2 → 2PbO It converts to lead sesquioxide on heating to 775k in air. 2PbO + O2 → 2Pb2O3 It is insoluble in water and is amphoteric. PbO + 2HNO3 → Pb(NO3)2 + H2O PbO + 2NaOH → Na2PbO2 + H2O Lead dioxide is obtained by the oxidation of lead acetate with bleaching powder in presence of slaked lime or by the fusion of lead monoxide with potassium nitrate or chlorate. 18 2. Halides All the four tetrahalides are known for carbon, silicon, germanium and tin. They are covalent compound and they all have the tetrahedral structure. The tetrahalides of Lead are unstable and are decomposed in presence of water. The dihalides are obtained by the addition of a soluble lead salt the corresponding hydrogen halide. Pb(NO3)2 + 2HX → PbX2 + 2HNO3 (X = F, Cl, Br, I) These are ionic compounds and are soluble in water. 3. Hydrides A number of stable covalent hydrides of carbon are known which are divided into series, viz. aliphatic compounds including alkanes (CnH2n + 2), alkenes (CnH2n), alkynes (CnH2n-2) and aromatic compounds (benzene, phenol, etc.). These are better described in organic chemistry classes. Silicon forms a number of silicon hydrides, known as silanes with the general formula SinH2n+2. These are analoguous to hydrocarbons but are comparatively much less stable and only a few are known due to limited capacity of silicon to show catenation. Monosilane is the most common hydride of silicon and is prepared by the action of dilute acid on magnesium silicate in a flask containing hydrogen gas. Mg2Si + 4HCl → 2MgCl2 + SiH4↑ Germanium forms hydrides or germanes, with the general formula GenH2n+2 (n = 1 to 5). These are obtained by the treatment of GeO2 with NaBH4 in acid solution by heating GeO2 with LiAlH4. The germanes are more resistant to hydrolysis than silanes. Their boiling points are also comparatively higher and they are less volatile and less flammable. Tin forms two hydrides SnH4, monostannane and Sn2H6, distannane. Monostannane may be obtained, as a highly toxic gas, by the action of hydrochloric acid on magnesium stannide. It freezes at -140oC and boils at -52oC. It decomposes rapidly at 150oC to form tin mirror. It is used as a reducing agent. SnH4 → Sn + 2H2 Distannane is obtained by the hydrogenation of stannites in presence of boranes. It is comparatively less stable and decomposes easily. 19 4. Carbides, silicides, germides, stannides and plumbides Carbon forms binary compounds, with more electropositive elements, known as carbides. The carbides are classified into four categories: Ionic or saltlike carbides, Covalent carbides, Metallic or interstitial carbides and Borderline or iron type carbides The Ionic or Saltlike Carbides These carbides are formed by metals of groups 1, 2, 13 (except boron), cadmium, coinage metals and lanthanides. These are prepared by strongly heating the metal, its oxide or hydride with carbon, carbon monoxide or a hydrocarbon in an electric arc. Examples: From metal and carbon: 2Be + C → Be2C From metal oxide and carbon: CaO + 3C → CaC2 + CO From metal and carbon monoxide: 6Al + 3CO → Al4C3 + Al2O3 From metal and hydrocarbon: 4Li + C2H2 → Li2C2 + 2LiH From metal oxide and hydrocarbon: 2Ag2O + C2H2 → Ag2C2 + H2O These carbides are colourless crystalline ionic compounds containing metal ions in the interstices between the carbon anions. These are nonconducting in solid state and easily hydrolysed to yield hydrocarbons. These carbides are further classified depending upon the product of hydrolysis obtained as follows. (i) Acetylides The ionic carbides which yield acetylene on hydrolysis are known as acetylides and are considered acetylene derivatives. For example, BeC2, MgC2, CaC2, Al2(C2)3, etc. These carbides contain C22− CaC2 + 2H2O → Ca(OH) 2 + HC≡CH Some carbides such as Cu2C2, Ag2C2, etc., are not considered as true acetylides as they do not yield acetylene on hydrolysis. (ii) Methanides The ionic carbides which yield methane on hydrolysis are known as methanides and are considered methane derivatives. For example, Al4C3, Be2C, Mn3C, etc. These carbides contains C4– ions. Al4C3 + 12H2O → 4Al(OH)3 + 3CH4 (iii) Allylides The ionic carbides which yield allylene on hydrolysis are known as allylides and are considered as the allylene derivatives. Mg2C3 is the only allylide and is considered to contain C34− or (C=C=C)4-. Mg2C3 + 4H2O → 2Mg(OH)2 + H2C=C=CH2 (iv) Mixed Carbides The ionic carbides which yield a mixture of several hydrocarbons on hydrolysis are known as mixed carbides. Carbides of Th and U are the few examples of these carbides. 20 (b) Covalent Carbides These carbides are formed by combination with hydrogen, boron, silicon and elements of Group 16 and 17. However, except the SiC and B4C, the other covalent carbides are better considered hydrides, sulphides, oxides and halides accordingly. B4C and SiC are obtained by reduction of their oxides with carbon in an electric furnace. 2B2O3 + 7C → B4C + 6CO SiO2 + 3C → SiC + 2CO These carbides are extremely hard, thermally stable and chemically inert. Hence, these are not attack by water and acid. Silicon forms compounds with more electropositive elements such as alkali metals, alkaline earth metals (except Be) and many transition metals. These are known as silicides due to presence of silicon anions and are usually obtained by heating the metal with either silicon or with silica in presence of carbon. The structures of the metal silicides are diverse, Some examples of their solid state structural types are: isolated Si atoms (e.g. Mg2Si, Ca2Si), Si2-units (e.g. U3Si2), Si4-units (e.g. NaSi, KSi, CsSi), Sin- chains (e.g. CaSi), planar or puckered hexagonal networks of Si atoms (e.g.b-USi2, CaSi2) and 3- dimensional network of Si `atoms (e.g. SrSi2, a-USi2). Silicides are hard materials, but their melting points are generally lower than those of the metal carbides. Treatment of Mg2Si with dilute acids gives mixtures of silanes. The properties of some silicides make them useful as refractory materials (e.g. Fe3Si and CrSi2). Fe3Si is used in magnetic tapes and disks to increase their thermal stability. Germanium, tin and lead do not form solid state binary compounds with metals. In contrast, the formation of Zintl phases and Zintl ions, which contain clusters of group 14 metal atoms, is characteristic of these elements. Historically, Zintl phases have been produced by the reduction of Ge, Sn or Pb by Na in liquid NH3. The synthesis of [Sn5]2 typifies the preparation of Zintl ions, and the use of the encapsulating ligand crypt-222 to bind an alkali metal counter-ion has played a crucial role in the development of Zintl ion chemistry. Thus, salts such as [K(crypt-222)]2[Sn5] and [Na(crypt-222)]4[Sn9] can be isolated. Modern technology allows low-temperature X-ray diffraction studies of sensitive (e.g. thermally unstable) compounds. It is therefore possible to investigate salts such as [Li(NH3)4]4[Pb9]:NH3 and [Li(NH3)4]4[Sn9]:NH3 which are formed by the direct reaction of an excess of Pb or Sn in solutions of lithium in liquid NH3. 21 5. Silicates Silicates constitute about 95% of the earth’s crust, in the form of silicate minierals, aluminosilicate clays or silica. The various forms of silicates in the use are ceramics, glass, granite, bricks, cement, rocks, sands, etc. The silicates can be regarded as the metal derivations of silicic acids. Silicates are obtained by the fusion of an alkali metal oxide or carbonate with sand at about 1400oC. Most silicates are insoluble in water, while sodium silicate is soluble in water. The insolubility of silicates owes to the stronger Si-O bond which is broken only by a strong reagent such as HF. Classification of Silicates: Depending upon the type of linkage, silicates are classified as: (i) Orthosilicates Silicates containing discrete tetrahydra units, without sharing of any oxygen atoms are termed orthosilicates and are represented by the formula M2(II)(SiO4), where is a divalent cation such as Be, Mg, Mn, Fe or Zn or the formula M(IV)(SiO4), where M(IV) is a tetravalent cation such as Zr. Structure of orthosilicates (discrete 𝑆𝑖𝑂44− unit) In some orthosilicates, the general formula is M3(II)M2(III)[(SiO4)3], where M(II) is a six- coordinated divalent metal cation such as Mg, Ca or Fe (II) and M (IV) is an eight coordinated trivalent cation such as Al, Cr or Fe (IV). These are known as garnets and are used as red gemstones. (ii) Pyrosilticates The silicates in which two tetrahedral units are linked by sharing of one oxygen atom to form (Si2O7)6- units are known as pyrosilicates. These are named so due to their formation from orthophosphates by heating. The pyrosilicates are rare and only a few examples are known such as thortveitite, Sc2(Si2O7), and hemimorphite, Zn3(Si2O7).Zn(OH)2.H2O. Structure of pyrosilicates ((Si2O7)6- ion) (iii) CyclicSilicates: The silicates in which two oxygen atoms per tetrahedra are shared to form a cyclic structure are known as cyclic silicates. (iv) Chain Silicates: Some silicates are formed by sharing of two oxygen atoms per tetrahedra to form single or double chains and are known as chain silicates. 22 (v) Sheet Silicate: The silicates formed by sharing of three oxygen atoms of each SiO4 tetrahedral unit with the adjacent tetrahedral to form an infinite two-dimensional sheet are known as sheet silicates (vi) Three-Dimensional: Silicates These silicates are formed by sharing of all the four oxygen atoms of an SiO4 tetrahedra with the adjacent tetrahedral to form a three-dimensional lattice of formula SiO2, as in quartz, tridymite or crystobalite. Due to absence of any negative charge, the metal ions are not present. However, some of the Si4+ may be replaced by Al3+ to give isomorphous structures with negative charge which is balanced by larger metal ions such as Na+, K+, Ca2+, or Ba2+. These structures can be represented as MI[AlSi3O8] and MII [Al2Si2O8] and are found in feldspars, zeolites and ultramarines. 6. Glass Glass is a solid solution of silicates with a number of oxides such as Na2O, K2O, PbO, ZnO, MgO, CaO, BaO, Al2O3 and B2O3. The composition of glass is not fixed and hence no definite formula can be assigned. It is not a true solid and has no definite melting point. Thus, it becomes soft on heating at a certain temperature and with further rise in temperature turns to a slightly viscous fluid, which flows freely as a liquid. The softened glass can be blown and molded into any desired shape. Glass is manufactured using various raw materials such as silica of uniform size, carbonates of sodium, potassium, carbonate and barium, nitrates of sodium and potassium and oxides of heavy metals such as zinc and lead. Barium carbonate is added to obtain glass of high refractive index. Zinc oxide and boric acid decrease the coefficient of expansion while lead oxide increases the refractive power of glass. Addition of calcium phosphate and oxides of arsenic and antimony produces opalescent glass. 23 COMPARATIVE ACCOUNT OF COMPOUNDS OF GROUP 14 ELEMENTS 1. Solubility and Hydrolysis of Tetrahalides: Carbon tetrahalides are insoluble in water and are not hydrolysed with water, whereas the tetrahalides of other elements are hydrolysed in water. This is due to the presence of vacant d-orbitals in these elements, which enables the acceptance of electrons from the oxygen atoms of two water molecules. Thus, the covalency of the elements (Si, Ge and Sn) increases to form a coordination compound, which dissociates to eliminate two molecules of hydrogen chloride and again coordinate with two molecules of water. Thus, the process is repeated to form orthosilicic acid. On the other hand, carbon cannot exceed its covalency beyond four due to absence of vacant d- orbitals and hence is extremely stable. This also justifies the absence of complexes of tetrahalides of carbon while, the tetrahalides of other elements readily form complexes due to increase in their covalency from four to six. For example, [SiF6]2-, [GeF6]2- and [SnCl6]2- etc. 2. CO2 is a Gas While SiO2 is a Solid: CO2 is constituted by two doubly bonded oxygen atoms linked to the carbon atom. The similar size and energy of 2p-orbitals of C-atom and 2p-orbitals of O atom enables the effective lateral overlap of these orbitals for the formation of pi bonds between C and O atoms. Thus, CO2 exists as a discrete molecule and exists as a gas. While the difference in size and energy of 3p-orbitals of Si-atom and 2p-orbitals of O-atom hinders the lateral overlap of these orbitals. So Si=O is not formed and tetravalency of Si is satisfied by the linkage of each Si atom to the four oxygen atoms resulting in a three dimensional network. Hence, SiO 2 exists as a solid. 3. Basic Character of Oxides: CO2 and SiO2 are acidic, while GeO2 and SnO2 are amphoteric. Thus, the acidic character decreases and the basic character increases on moving down the group.