2024 Group 1, II, IVA Elements PDF

Group one-The Alkali metals Element Symbol Electronic Structure Lithium Li 1s22s1 Sodium Na 1s22s22p63s1 Potassium K 1s22s22p63s23p64s1 Rubidium Rb 1s22s22p63s23p64s24P65s1 Caesium Cs 1s22s22p63s23p64s24P65s15P66s1 Francium Fr [Rn]7s1  Occurrence and Abundance Despite t...

Group one-The Alkali metals Element Symbol Electronic Structure Lithium Li 1s22s1 Sodium Na 1s22s22p63s1 Potassium K 1s22s22p63s23p64s1 Rubidium Rb 1s22s22p63s23p64s24P65s1 Caesium Cs 1s22s22p63s23p64s24P65s15P66s1 Francium Fr [Rn]7s1  Occurrence and Abundance Despite their close chemical similarity, the elements do not occur together, mainly because their ions are of different size. Sodium is the most abundant, followed by potassium, rubidium, lithium, and caesium. Francium is intensely radioactive and very rare.  Lithium is the thirty-fifth most abundant element by weight and is mainly obtained as the silicate minerals, spodumene LiAl(SiO3)2 and lepidolite Li2Al2(SiO3)3(FOH)2.  Sodium and potassium are the seventh and eighth most abundant elements by weight in the earth’s crust. The largest source of sodium is rock salt (NaCl). Various salts including NaCl, Na2B4O7.10H2O (borax), NaNO3 (saltpetre) and Na2SO4 (mirabilite) are obtained from deposits formed by the evaporation of ancient seas.  Potassium occurs mainly as deposits of KCl (sylvite), a mixture of KCl and NaCl (Sylvinite), and the double salt KCl.MgCl2.6H2O (carnalite).  There is no convenient source of rubidium and only one of caesium and these elements are obtained as by-products from lithium processing.  All of the elements heavier than bismuth (atomic number 83) 83 Bi are radioactive. Thus francium (atomic number 89) is radioactive and has a short half-life period of 21 minutes it does not occur appreciably in nature.  Extraction of metals The metals may all be isolated by electrolysis of a fused salt, usually the fused halide, often with impurity added to lower the melting point. Sodium is made by the electrolysis of a molten mixture of about 40% NaCl and 60%CaCl2 in a Downs cell.This mixture melts at about 600C compared with 803C for pure NaCl. The small amount of calcium formed during the electrolysis is insoluble in the liquid sodium, and dissolves in the eutectic mixture. Eutectic mixture is a type of a homogeneous mixture that has a melting point lower than those of the constituents.  There are three advantages to electrolyzing It lowers the melting point and so reduces the fuel bill. The lower temperature results in a lower vapour pressure for sodium, which is important as sodium vapour ignites in air.  At lower temperature the liberated sodium metal does not dissolve in the melt, and this is important  because if it dissolved it would short-circuit the electrodes and thus prevent further electrolysis. A Downs cell comprises a cylindrical steel vessel lined with firebrick, measuring about 2.5m in height and 1.5m in diameter.  The anode is a graphite rod in the middle, and is surrounded by a cast steel cathode. A metal gauze screen separates the two electrodes, and prevents the Na formed at the cathode from recombining with Cl2 produced at the anode. The molten sodium rises, as it is less dense than the electrolyte, and it is collected in an inverted trough  and removed, and packed into steel drums. A similar cell can be used to obtain potassium by electrolyzing fused KCl. However, the cell must be operated at a higher temperature because the melting point of KCl is higher, and this results in the  vapourization of the liberated potassium. Since sodium is readily available, the modern method is to reduce molten KCl with sodium vapour at 850 C in a large fractionating tower. This gives K of 99.5% purity. Na +KCl →NaCl +K Rb and Cs are produced in a similar way by reducing the chlorides with Ca at 750C under reduced pressure. All react with water to give hydrogen gas and the metal hydroxide. They also react with the oxygen in the air to give either an oxide, peroxide or superoxide depending on the metal. These metals almost always form ions with a positive (+1) charge.  Most of the alkali metals glow with a characteristic colour when placed in a flame; lithium is crimson, sodium gives off an intense yellow, potassium is Lilac, rubidium is a red-violet, and caesium gives off blue light.  Electronic structure Group 1 elements all have one valence electron in their outer orbital- an s electron, which occupies a spherical orbital. The single valence electron is a long distance from the nucleus, is only weakly held and is readily removed. In contrast the remaining electrons are closer to the nucleus, more tightly held, and are removed only with great difficulty. Because of similarities in the electronic structures of these elements, many similarities in chemical behaviour would be expected.  Size of atoms and ions Group 1 atoms are the largest in their periods in the periodic table. When the outer electron is removed to give a positive ion, the size decreases considerably. There are two reasons for this. 1). The outermost shell of electrons has been completely removed. 2). Having removed an electron, the positive charge on the nucleus is now greater than the charge on the remaining  electrons, so that each of the remaining electrons is attracted more strongly towards the nucleus. This reduces the size further.  Positive ions are always smaller than the parent atom. Even so, the ions are very large, and they increase in size from Li+ to Fr+ as extra shells of electrons are added. The Li is much smaller than + the other ions. For this reason, Li only mixes with Na above 380°C and it is immiscible with the metals K, Rb and Cs, even when molten. In contrast the other metals Na, K, Rb and Cs are miscible with each other in all proportions. Density The atoms are large, so group 1 elements have remarkably low densities.  Metalli Ionic Density c radius (gcm-3) radius M+(Å) (Å) Li 1.52 0.76 0.54 Na 1.86 1.02 0.97 K 2.27 1.38 0.86 Rb 2.48 1.52 1.53 Cs 2.65 1.67 1.90 Ionization Energy The first ionization energies for the atoms in this group are appreciably lower than those for any other group in the periodic table. The atoms are very large so the outer electrons are only held weakly by the nucleus.  Hence the amount of energy needed to remove the outer electron is not very large.  On descending the group from Li to Na to K to Rb to Cs, the size of the atoms increases; the outermost electrons become less strongly held, so the ionization energy decreases. Electronegativity The electronegativity values for the elements in this group  are very small- in fact the smallest values of any element. Thus when these elements react with other elements to form compounds, a large electronegativity difference between the two atoms is probable, and ionic bonds are formed. Li- 1.0, Na-0.9, K-0.8, Rb-0.8, Cs-0.7(Pauling’s electronegativity). Chemical properties Some reactions of Group 1 metals 2M +2H2O →2MOH +H2 The hydroxides are the strongest bases known. With excess dioxygen 4Li +O2 →2Li2O Monoxide is formed by Li and to a small extent by Na.  2Na+O2→Na2O2 Peroxide is formed by Na and to a small extent by Li K +O2 →KO2 Superoxide  2M +H2 →2MH ionic ‘salt-like’ hydrides. 6Li +N2 →2Li3N Nitride formed only by Li. 3M +P →M3P All the metals form phosphides  3M +As →M3As All the metals form arsenides 3M +Sb →M3Sb All the metals form stibnides  2M +X →M2X (X=S,Se,Te) All the metals form sulphides, selenides, and tellurides.  2M +X2 →2MX (X=F, Cl, Br, I) All the metals form fluorides, chlorides, bromides, and iodides 2M + 2NH3 →2MNH2 + H2 All the metals form amides  Uses of Lithium: Lithium is used to make electrochemical cell (both primary and secondary batteries). Lithium is used in lubricants, in glass industries, and in alloys of lead, aluminum, and magnesium to make them less dense and stronger. Uses of sodium: Liquid sodium metal is used as a coolant in fast breeder nuclear reactor. Sodium has many biological uses like nerve signal transmission. Sodium nitrite is a principal ingredient in gunpowder. The pulp and paper industry uses large amounts of sodium hydroxide, sodium carbonate, and sodium sulphate.  Sodium carbonate is used by power companies to absorb sulfur dioxide, a serious pollutant, from smokestack gases (locomotive chimneys or  ship chimneys). Sodium carbonate is also used in the glass and detergent industries. Sodium chloride is used in foods and to soften the water. Sodium bicarbonate (baking soda) is used in the food industry as well.  Uses of potassium: Potassium is an essential element for life. Roughly 95% of Potassium compounds are used as fertilizers for plants.  Potassium hydroxide is used in detergent. Potassium chlorate is used in explosive.  Potassium carbonate is used in ceramics, colour TV tubes and fluorescent light tubes. Potassium bromide is used in photography industries. Uses of Rubidium and caesium: Rubidium is used almost exclusively for research, but caesium is used in special glasses and radiation detection equipment GROUP TWO-THE ALKALINE EARTH METALS Element Symbol Electronic Structure Beryllium Be 1s22s2 Magnesium Mg 1s22s22p63s2 Calcium Ca 1s22s22p63s23p64s2 Strontium Sr 1s22s22p63s23p63d104s24p65s2 Barium Ba 1s22s22p63s23p63d104s24p64d105s25p66s2 Radium Ra [Rn]7s2  Alkaline earth metals make up the second group of the periodic table. This family includes the elements beryllium, magnesium, calcium, strontium, barium, and radium (Be, Mg, Ca, Sr, Ba, and Ra, respectively). These metals are silver and soft, much like the alkali metals of Group 1.  Each alkaline earth metal has two valence electrons. They will easily give these electrons up to form cations. These metals become increasingly more reactive as you go down the periodic table. This is concurrent with general periodic trends.  The group two elements show the same trends in properties as were observed with Group 1. However, beryllium stands apart from the rest of the group and differs much more from them than lithium does from the rest of Group 1.  The main reason for this is that the beryllium atom and Be2+ are both extremely small, and the relative increase in size from Be2+ to Mg2+ is four times greater than the increase between Li+ and Na+.  Beryllium and barium compounds are all very toxic. The elements form a well-graded series of highly reactive metals, but are less reactive than Group 1. They are typically divalent and generally form colourless ionic compounds.  The oxides and hydroxides are less basic than those of Group 1: hence their oxosalts (carbonates, sulphates, nitrates) are less stable to heat. Occurrence and Extraction These elements are all found in the Earth’s crust, but not in the elemental form as they are so reactive. Instead, they are widely distributed in rock structures.  Beryllium, like its neighbours Li and B is relatively not very abundant in the earth’s crust. It occurs to the extent of about 2ppm and is thus similar to Sn (2.1 ppm), Eu (2.1 ppm) and As (1.8 ppm). Beryllium is found in small quantities as the silicate minerals beryl Be3Al2Si6O18 and phenacite Be2SiO4.  Magnesium is the sixth most abundant element in the earth’s crust (27640 ppm ). The main minerals in which magnesium is found are carnellite (KCl.MgCl2.6H2O), magnesite (MgCO3) and dolomite (MgCO3.CaCO3).  Calcium is the fifth most abundant element in the earth’s crust. Hence the third most abundant metal after Al and Fe. Calcium is found in gypsum (CaSO4.2H2O), anhydrite (CaSO4), fluorite (CaF2), apatite (Ca5(PO4)3F) and limestone (CaCO3) There are two crystalline forms of CaCO3, calcite and aragonite.  Strontium (384ppm) and barium (390ppm) are much less abundant, but are well known because they occur as concentrated ores, which are easy to extract. They are respectively the fifteenth and fourteen element in abundance.  Strontium is mined as celestite SrSO4 and strontianite (SrCO3). Ba is mined as Barytes BaSO4. Radium is extremely scarce and is radioactive. It was first isolated by Pierre and Marie Curie by processing many tons of the uranium ore pitchblende.  Of the elements in this Group only magnesium is produced on a large scale. It is extracted from sea-water by the addition of calcium hydroxide, which precipitates out the less soluble magnesium hydroxide. This hydroxide is then converted to the chloride, which is electrolysed in a Downs cell to extract  The metals of this group are not easy to produce by chemical reduction because they are themselves strong reducing agents, and they react with carbon to form carbides. They are strongly electropositive and react with water, and so aqueous solutions cannot be used for displacing them with another metal, or for electrolytic production.  All the metals can be obtained by electrolysis of the fused chloride, with sodium chloride added to lower the melting point, although strontium and barium tend to form a colloidal suspension. Properties of the elements  The alkaline earth metals are silvery white, lustrous and relatively soft. Their physical properties when compared with those of group 1A metals, show that they have a substantially higher melting point., boiling point, enthalpies of fusion and vapourization.  They have two valency electrons which may participate in metallic bonding, compared with one electron for Group 1 Metals. Consequently Group 2 metals are harder, have higher cohesive energy and much higher melting points and boiling points than Group 1 elements.  The melting points do not vary regularly, mainly because the metals adopt different crystal structures. Melting points of Group 1 and 2 Melting Pt(°C) Be 1287 Li 181 Mg 649 Na 98 Ca 839 K 6 Sr 768 Rb 39 Ba 727 Cs 28.5  This can be understood in terms of the size factor and the fact that two valency electrons per atom are now available for bonding. Again, Be is notable in melting more than 1100°C above Li and being nearly 3.5 times as dense; its enthalpy of fusion is more than 5times that of Li.  Size of atoms and ions Group 2 atoms are large but are smaller than the corresponding group 1 elements as the extra charge on the nucleus draws the orbital electrons in. Similarly the ions are large, but smaller than those of Group 1, especially because of the removal of two orbital electrons increases the effective nuclear charge further.  Thus, these elements have higher densities than group 1 metals.  Size and Density metallic Ionic Density radius(Å) Radius(Å) (gcm-3) Be 1.12 0.31 1.85 Mg 1.60 0.72 1.74 Ca 1.97 1.00 1.55 Sr 2.15 1.18 2.63 Ba 2.22 1.35 3.62 Ra 1.48 5.5  Comparison with group 1A shows the substantial increase in the ionization energies; this is related to their smaller size and higher nuclear charge and is particularly notable for Be. Ionization energy 1st 2nd 3rd Be 899 1757 14847 Mg 737 1450 7731 Ca 590 1145 4910 Sr 549 1064 Ba 503 765 Ra 509 979  The third ionization energy is so high that M3+ ions are never formed. The ionization energy for Be2+ is high and its compounds are typically covalent, Mg also forms some covalent compounds. However, the compounds formed by Mg, Ca, Sr and Ba are predominantly divalent and ionic.  Since the atoms are smaller than those in Group 1, the electrons are more tightly held so that the energy needed to remove the first electron (1st ionization energy) is greater than those in Group 1.  Once one electron has been removed, the ratio of charges on the nucleus to orbital electrons is increased, so that the remaining electrons are more tightly held. Hence the energy needed to remove a second electron is nearly double that required for the first. ELECTRONEGATIVITY The electronegativity values of Group 2 elements are low, but are higher than the values for Group 1. The electronegativity difference between Group 2 metals (Mg, Ca, Sr, and Ba) and the halogens or oxygen is large and the compounds are ionic. The value for Beryllium is higher than for others. HYDRATION ENERGIES The hydration energies of the Group 2 ions are four or five times greater than for group 1 ions. This is largely due to their smaller size and increased charge. The crystalline compounds of Group 2 contain more water of crystallization than the corresponding Group 1 compounds.  Thus NaCl and KCl are anhydrous but MgCl2.6H2O, CaCl2.6H2O and BaCl2.2H2O all have water of crystallization. Note that the number of molecules of water of crystallization decreases as the ions become larger. Differences between Beryllium and the other Group 2 elements Beryllium is anomalous in many of its properties and shows diagonal relationship to aluminum in Group 3. It is extremely small and has a high charge density and so by Fajans rules it has a strong tendency to covalency.  Beryllium hydride is electron deficient and polymeric with multicentre bonding like aluminium hydride. The halides of beryllium are electron deficient, and polymeric with halogen bridges. BeCl2 usually forms chains but also exists as the dimer. AlCl3 is dimeric. Be forms many complexes – not typical of Groups 1 and 2.  Be like Al is rendered passive by nitric acid. Be is amphoteric, liberating H2 with NaOH and forming beryllates. Al forms Aluminates. Be(OH)2 like Al(OH)3 is amphoteric. Be salts are extensively hydrolysed. Be salts are among the most soluble known.  Beryllium forms an unusual carbide Be2C which like Al4C3 reacts with water to give methane whereas magnesium carbide and calcium carbide give propyne and ethyne(formerly called acetylene) respectively. Be2C+4H2O → 2Be(OH)2 + CH4 Mg2C3 + 4H2O → 2Mg(OH)2 + C3H4  Beryllium metal is relatively unreactive at room temperature, particularly in its massive form. It does not react with water or steam even at red heat and does not oxidize in air below 600°, though powdered Be burns brilliantly on ignition to give BeO and Be3N2.  The halogens (X2) react above 600°C to give BeX2 but the chalcogens (S, Se, Te) require higher temperatures to form BeS, e.t.c. Ammonia reacts above 1200°C to give Be3N2 and carbon forms Be2C at 1700°C. In contrast with the other group IIA  Cold concentrated HNO3 passivates Be but the metal dissolves readily in dilute aqueous acids (HCl, H2SO4, HNO3) with the evolution of hydrogen. Beryllium is sharply distinguished from the other alkaline earth metals in reacting with aqueous alkalis(NaOH, KOH) with evolution of hydrogen.  Magnesium is more electropositive than the amphoteric Be and reacts more readily with most of the non metals. It ignites with the halogens, particularly when they are moist, to give MgX2 and burns with dazzling brilliance in air to give MgO, Mg3N2.  It also reacts directly with the other elements in Group V and VI (and Group IV); when heated and even forms MgH2 with hydrogen at 570 and 200 atm. Steam produces MgO, or Mg(OH)2 plus Hydrogen and ammonia reacts at elevated temperature to give Mg3N2.  The heavier alkaline earth metals, Ca, Sr, Ba (and Ra) react even more readily with non metals and again the direct formation of nitrides M3N2 is notable. Other products are similar though the hydrides are more stable and the carbides less stable than for Be and Mg.  There is also a tendency, previously noted for the alkali metals to form peroxides MO2 of increasing stability in addition to the normal oxides MO.  Some reactions of Group 2 metals M +2H2O →M(OH)2 +H2 Mg with hot water, and Ca, Sr and Ba react rapidly with cold water. 2M +O2 →2MO Normal oxide formed by all group members. With excess dioxygen Ba +O2 →BaO2 Ba also forms the peroxide. M +H2 →MH2 Ionic ‘salt–like’ hydrides formed at high temperatures by Ca, Sr and Ba. 3M +N2 →M3N2 All form nitrides at high temperatures.  3M +2P →M3P2 All the metals form phosphides at high temperatures. M +X →MX (X=S,Se,Te) All the metals form sulphides, selenides, and tellurides. M +X2 →MX2 (X=F, Cl,Br, I) All the metals form fluorides, chlorides, bromides and iodides.  M +2NH3 →M(NH2)2 + H2 All the metals form amides at high temperatures. Hydroxides Be(OH)2 is amphoteric, but the hydroxides of Mg, Ca, Sr and Ba are basic. The basic strength increases from Mg to Ba and Group 2 shows the usual trend that basic properties increase on descending a group. SULPHATES The solubility of the sulphates in water decreases down the group, Be  Mg  Ca  Sr  Ba. Thus BeSO4 and MgSO4 are soluble but CaSO4 is sparingly soluble and the sulphates of Sr, Ba and Ra are virtually insoluble.  The significantly higher solubilities of BeSO4 and MgSO4 are due to the high enthalpy of solvation of the smaller Be2+ and Mg2+ ions. The sulphates all decompose on heating, giving the oxides: heat MgSO4 → MgO+SO3  MgSO4 and CaSO4 cause permanent hardness in water while the presence of Mg(HCO3)2 and Ca(HCO3)2 causes temporary hardness in water.  HYDRIDES The elements Mg, Ca, Sr and Ba all react with hydrogen to form hydrides MH2. Beryllium hydride is difficult to prepare, and less stable than the others. Hydrides are all reducing agents and are hydrolysed by water and dilute acids with the evolution of hydrogen.  CaH2 +2H2O→Ca(OH)2+2H2 NITRIDES The alkaline earth elements all burn in dinitrogen and form ionic nitrides M3N2. This is in contrast to Group 1 where Li3N is the only nitride formed. 3Ca +N2 →Ca3N2 All the nitrides are all crystalline solids, which decompose on heating and react with water, liberating ammonia and forming either the metal oxide or hydroxide e.g. Ca3N2 +6H2O →3Ca(OH)2 +2NH3  Compounds The predominant divalence of the Group IIA metals can be interpreted in terms of their electronic configuration, ionization energies and size. Further ionization to MX3 is impossible[14847kJmol-1 for Be, 7731kJmol-1 for Mg and 4910kJmol-1 for calcium].  USES In its elemental form, magnesium is used for structural purposes in car engines, pencil sharpeners, and many electronic devices such as laptops and cell phones. In a biological sense, magnesium is vital to the body's health: the Mg2+ ion is a component of every cell type.  Calcium metal is used to make alloys with Aluminium for bearings. It is used in the iron and steel industry to control carbon in cast iron and as a scavenger for P, O and S. Other uses are as a reducing agent in the production of other metals-Zr, Cr, Th and U- and for Chemically CaH2 is sometimes used as a source of H2. Radium was used for radiotherapy treatment of cancer at one time: other forms of radiation are now used.  GROUP 4A ELEMENTS Element Symbol Electronic Structure Carbon C [He] 2s22p2 Silicon Si [Ne] 3s23p2 Germanium Ge [Ar] 3d104s24p2 Tin Sn [Kr] 4d105s25p2 Lead Pb [Xe] 4f145d106s26p2  OCCURRENCE OF THE ELEMENTS The elements are all well known, apart from germanium. Carbon is the seventeenth and silicon the second most abundant element by weight in the earth’s crust. Germanium minerals are very rare. Germanium occurs as traces in the  other metals and in coal, but it is not well known. The abundances of tin and lead are comparatively low, they occur as concentrated ores which are easy to extract.  Carbon occurs in large quantities combined with other elements and compounds mainly as coal, crude oil, and carbonates in rocks such as calcite CaCO3, magnesite MgCO3 and dolomite [MgCO3. CaCO3].  Silicon occurs very widely, as silica SiO2 (sand and quartz) and in a wide variety of silicate minerals and clays. Germanium is only found as traces in some silver and zinc ores. Tin is mined as cassiterite SnO2. lead is found as the ore galena PbS.  Extraction Carbon black (soot) is produced in large amounts. It is made by the incomplete combustion of hydrocarbons from natural gas or oil. Natural graphite is usually found as a mixture with mica, quartz and silicates, which contains 10-60% C.  Silicon is made by reducing SiO2 and scrap iron with coke. SiO2 + Fe + 2C FeSi + 2CO There must be an excess of SiO2 to prevent the formation of the carbide SiC. High purity silicon is made by converting Si to SiCl4, purifying this by distillation, and reducing the chloride with Mg or Zn.  SiO2 + 2C Si + 2CO Si + 2Cl2 SiCl4 SiCl4 + 2Mg Si + 2MgCl2 The only important ore of tin is SnO2(Cassiterite). SnO2 is reduced to the metal using Carbon at 1200-1300C in an electric furnace. The product often contains traces of iron, which make the metal hard.  Iron is removed by blowing air through the molten mixture to oxidize the iron to FeO, which then floats to the surface. The main oxide of lead is galena PbS. There are two methods of extracting the element: 1). Roast in air to give PbO and then reduce with coke or CO in a blast furnace. 2PbS + 3O2 2PbO + 2SO2 +C 2Pb(liquid) + CO2(gas)  2). PbS is partially oxidized by heating and blowing air through it. The air is turned off after some time, heating continues and the mixture undergoes self-reduction. 3PbS heat in heat in PbS + 2PbO 3Pb(liquid) + air air SO2(gas)  The Group 4A elements are found in the p-block. Each of these elements has only two electrons in its outermost p orbital with electron configuration ns2np2. The Group 4A elements tend to adopt oxidation states of +4 and, for the heavier elements, +2 due to the inert pair effect.  The inert pair effect is the phenomenon of electrons remaining paired in valence shell. It can be defined as the reluctance of the outermost shell s-electrons to participate in bonding.  Members of this group conform well to general periodic trends. The atomic radii increase down the group, and ionization energies decrease.  Metallic properties increase down the group. Carbon and silicon are non- metals, germanium has some metallic properties, tin and lead are metals.  The elements in this group are relatively unreactive, but reactivity increases down the group. MII oxidation state becomes increasingly stable on descending the group. Pb often appears more noble(unreactive) than expected from its standard electrode potential of - 0.13volts.  The unreactiveness is partly due to a surface coating of oxide. Carbon, silicon and germanium are unaffected by water, tin reacts with steam to give SnO2 and H2. Pb is unaffected by water, probably because of a protective oxide film.  Carbon is the fourth most abundant element in the known universe but not nearly as common on the earth, despite the fact that living organisms contain significant amounts of the element.  Common carbon compounds in the environment include the gases carbon dioxide (CO2) and methane (CH4). Allotropes Carbon exists in several forms called allotropes. Diamond is one form with a very strong crystal lattice, known as a precious gem from the most ancient records.  Graphite is another allotrope in which the carbon atoms are arranged in planes which are loosely attracted to one another (hence its use as a lubricant). The recently discovered fullerenes are yet another form of carbon.  Inorganic carbon may come in the form of diamond as transparent, isotropic crystal. It is the hardest natural occurring material on this earth. Diamond has four valence electrons, and when each electron bonds with another carbon it creates a sp3-hybridized atom. The boiling point of diamond is 4827°C.  Unlike diamond, graphite is opaque, soft, dull and hexagonal. Graphite can be used as a conductor (electrodes) or even as pencils.  Germanium, categorized as a metalloid in group 4A, the Carbon family, has five naturally occurring isotopes. Germanium, abundant in the Earth's crust has been said to improve the immune system of cancer patients. It is also used in transistors, but its most important use is in fiber-optic systems and infrared optics.  The name for silicon is taken from the Latin silex which means "flint". The element is second only to oxygen in abundance in the earth's crust and was discovered by Berzelius in 1824. The most common compound of silicon, SiO2, is the most abundant chemical compound in the earth's crust, which we know it better as common beach sand.  Properties Silicon is a crystalline semi- metal or metalloid. One of its forms is shiny, grey and very brittle (it will shatter when struck with a hammer). It is a group 4A element in the same periodic group as carbon, but chemically behaves distinctly from all of its group counterparts.  Silicon shares the bonding versatility of carbon, with its four valence electrons, but is otherwise a relatively inert element. However, under special conditions, silicon can be made to be a good deal more reactive.  Silicon exhibits metalloid properties, is able to expand its valence shell, and is able to be transformed into a semiconductor; distinguishing it from its periodic group members.  27.6% of the Earth's crust is made up of silicon. Although it is so abundant, it is not usually found in its pure state, but rather its dioxide and hydrates. SiO2 is silicon's only stable oxide, and is found in many crystalline varieties.  Its purest form being quartz, but also as jasper and opal. Silicon can also be found in feldspar, micas, olivines, pyroxenes and even in water. Silicon is most commonly found in silicate compounds.  Applications Carbon has a very high melting and boiling point and rapidly combines with oxygen at elevated temperatures. In small amounts it is an excellent hardener for iron, yielding the various steel alloys upon which so much of modern construction depends.  Activated carbon is used extensively in sugar industry as decolourizing agent. It is also used in purification of chemicals and gases. It is used as catalyst.  An important (but rare) radioactive isotope of carbon, C-14, is used to date ancient objects of organic origin. It has a half-life of 5730 years but there is only 1 atom of C-14 for every 1012 atoms of C-12 (the usual isotope of carbon).  Silicon is a semiconductor with a clear shiny bluish grey metallic lustre. It is used in the production of transistors while an isotope 29Si, is used in NMR spectroscopic studies.  Si/steel alloys are used for the construction of electric motors. Silicon dioxide and Silicon (in the form of clay or sand) are important components of bricks, concrete and Portland cement. Silicon  Silicon parts are used in computers, transistors, solar cells, Liquids Crystal Display (LCD) screens and other semiconductors devices. Silicates are used to make pottery and enamel. Sand which contains silicon is an important component of glass. Silicones are used in high temperature greases, waxes, breast implants, contact lenses, explosives and pyrotechnics.  Germanium is transparent to infra-red light and is therefore used for making prisms and lenses and windows in infra-red spectrophotometer.  Germanium is used in transistor technology and in optics. Magnesium germanate is used in the special alloy, strain gauges and in superconductors. Germanium is used in electronic application when doped with arsenic, gallium or other elements.  Germanium oxide has a high refractive index of refraction and dispersion, making it suitable for use in wide-angle camera lenses and objectives lenses for microscopes. It is also used as an alloying agent (adding 1% germanium to silver stops if from tarnishing) in fluorescent lamps and so a catalyst.  Both Ge and germanium oxide are transparent to infrared radiation and so are used in the manufacture of infrared spectroscopes.  Tin is used for electroplating steel to make tin-plate and alloys. Lead is used to make lead/acid storage batteries. Lead is used as protective shielding against X-ray and radiation from nuclear reactors