Inorganic Chemistry

INORGANIC CHEMISTRY – CLASS XI (ISC)
Properties of Group 1 elements [Alkali metals] 1. Due to high reactivity, alkali metals do not occur free in nature. Elements of group 1 (or IA) are known as alkali metals because their hydroxides are soluble in water and form strongly alkaline solutions. Alkali metals are stored under kerosene oil because they get tarnished on exposure to air. 2. The general electron configuration of alkali metals is ns1. 3. Alkali metals have largest size and lowest ionization enthalpy in their respective periods. Down the group, the size of alkali metals increases whereas ionization enthalpy decreases. 4. Alkali metals have low densities and densities increases from Li to Cs. (Exception: K is lighter than Na, due to bigger size of potassium atom). 5. Alkali metals have low melting and boiling points due to their loosely packed metallic lattices. 6. Alkali metals, except lithium, show photoelectric effect. Due to small size and high ionization enthalpy, the electrons of lithium cannot be emitted. 7. Alkali metals do not form dipositive ions because of their very high second ionization enthalpies. 8. All alkali metals are paramagnetic but their salts are diamagnetic. 9. Alkali metals impart characteristic colors to the flame, i.e., lithium imparts crimson red, sodium imparts golden yellow, potassium imparts pale violet while rubidium and caesium impart violet color to the flame. This is due to the fact that the atoms absorb energy from the flame and the electrons jump from lower orbit to higher orbit. However, since ionization enthalpies of the lower atoms of the group are higher, the jump is small and the energy (in the form of visible light) radiated when the electrons come back to their original positions were also small (i.e. with low frequencies). Since red color is of low frequency, we see red color for the lowest element of the group, i.e. lithium. Thereafter, the frequency of the emitted light for the higher elements gradually increases. 10. Lithium burns in oxygen to form lithium monoxide while the other alkali metals form peroxides. Li does not form peroxides because the strong positive field around the lithium ion (Li+) attracts the negative charge so strongly that it does not permit monoxide (O–) ion to combine with oxygen to form peroxide ion (O22–). 11. The peroxide anions are diamagnetic and are oxidising agents. 12. Na2O2 reacts with CO2 of the air and has been used in submarines and confined places, as it removes both CO2 and produces O2. 2Na2O2 + 2CO2 → 2Na2CO3 + O2 13. Superoxides contain the ion [O2–] and are paramagnetic and colored due to the presence of an unpaired electron. LiO2 and NaO2 are yellow, KO2 is orange. RbO2 is brown and CsO2 is orange.
Superoxides are stronger oxidising agents than peroxides and give H2O2 and O2 on treatment with water or acids. Potassium superoxide [KO2] is used in breathing masks, submarines and space capsules because it removes CO2 and at the same time produces O2. 4KO2 + 2CO2 → 2K2CO3 + 3O2 4KO2 + 4CO2 + 2H2O → 4KHCO3 + 3O2 14. Lithium cannot be stored in kerosene oil as it floats on the surface of oil due to low density. Therefore, it is kept wrapped in paraffin wax. All alkali metals dissolve in mercury forming amalgams with evolution of heat. Li is the exception. It is used as scavenger in the metallurgical operations to remove oxygen and nitrogen. It is the lightest metal known. It is least fusible, least dense and least soft of all alkali metals. It has highest specific heat among all the alkali metals and specific heat decreases from Li to Cs. 15. Due to large negative electrode potentials, alkali metals are strong reducing agents. The reducing character increases from Na to Cs. However, lithium is the strongest reducing agent among all the alkali metals in spite of its highest Ionization enthalpy (IE). This is because of extensive hydration of Li+ ions and large amount of energy released during hydration, more than compensates the higher IE value of lithium. Alkali metals are better reducing agents than H2. Therefore, they react with compounds containing acidic hydrogen atoms liberating H2 gas. 16. Alkali metal halides are ionic in nature and have high melting points due to their ionic nature. The melting points follow the order: Fluorides > chlorides > Bromides > Iodides. This is due to progressive decrease in ionic character. Melting points of chlorides follow the order : LiCl < NaCl > KCl > RbCl > CsCl. Low melting point of LiCl as compared to NaCl is most probably due to the covalent nature of LiCl. Covalent character of lithium halides follows the order: LiI > LiBr > LiCl > LiF. 17. Solutions of alkali metals in liquid ammonia are highly conducting and deep blue in color. This is because of ammoniated cations and electrons. Solution is paramagnetic due to the presence of a large number of unpaired electrons. 18. Unlike other members of the group, lithium forms binary compound with nitrogen, i.e. Li3N.

Properties of Group 2 elements [Alkaline earth metals] 1. Elements of group 2 (or IIA) are called alkaline earth, since: (i) Their oxides are alkaline in nature like alkali metal oxides, and, (ii) the oxides of Ca, Sr and Ba are found in earth's surface. The general electron configuration of alkaline earth metals is ns2. The last alkaline earth metal, i.e., Ra is radioactive. 2. Due to higher nuclear charge, the atomic and ionic radii of alkaline earth metals are smaller than their corresponding alkali metals. 3. Bivalent positive ions of alkaline earth metals are stable than M+ ions in spite of the fact that IE2 of alkaline earth metals is almost double than the IE1. This is due to their high lattice energies in the solid state and high hydration energies in aqueous solutions. 4. Alkaline earth metals have nearly zero values for electron affinity (electron gain enthalpy) due to their stable configurations. 5. Like alkali metals, alkaline earth metals (except Be and Mg) also impart characteristic colors to the flame, e.g., Ca imparts brick red, strontium gives crimson and barium gives green color. Be and Mg do not impart any characteristic color to the flame because of their high IE. 6. Alkaline earth metals are harder and denser than alkali metals due to their more close packed structure. 7. The alkaline earth metals are weaker reducing agents than alkali metals since their standard electrode (reduction) potentials are less negative than their corresponding alkali metals. 8. Due to their higher ionization enthalpies and less solubility, the alkaline earth metal hydroxides are weaker than corresponding alkali metal hydroxides. Basic strength of the hydroxides decreases down the group. Thus, basic strength follows the order: Be(OH)2 < Mg(OH)2 < Ca (OH)2 < Sr(OH)2 < Ba (OH). Be(OH)2 is, however, amphoteric. 9. Alkaline earth metals and their divalent ions are colorless and diamagnetic due to the absence of unpaired electrons. 10. In general, dissolution of a salt takes place when hydration energy is greater than the lattice energy of the ionic solid. Thus, solubility decreases in the order: BeSO4 > MgSO4 > CaSO4 > SrSO4 > BaSO4. 11. Beryllium halides are covalent, whereas other alkaline earth metal halides are ionic. Beryllium halides are soluble in organic solvents whereas the halides (except fluorides) of other alkaline earth metals are soluble in water. Fluorides, however, are insoluble in water due to large values of lattice energy. 12. Be2C, on reaction with water, gives methane whereas carbides of other alkaline earth metals give acetylene gas.

Properties of Group 13 elements 1. The atomic and ionic radii of group 13 elements are smaller than the corresponding elements of group 2.This is because on moving from left to right, i.e., from group 2 to group 13 in a given period, the nuclear charge increases while the new electron enters the same shell. Further the electrons in the same shell do not screen each other. Therefore, the effective nuclear charge increases and the electrons are pulled more towards the nucleus. This results in decrease in atomic size. Same is true of ionic radius. On moving down the group, both atomic and ionic radii are expected to increase primarily due to addition of a new electron shell with each succeeding element. 2. The atomic radius of Ga (135 pm) is slightly lower than that of Al (143 pm).This is due to the filling of electrons in d-orbitals. In-between Al (Z = 13) and Ga (Z = 31), there are ten elements of the first transition series (Z = 21 to 30) which have electrons in the inner d-orbitals. As the d-orbitals are large in size, these intervening electrons do not screen the nucleus effectively. Consequently, effective nuclear charge of Ga is greater in magnitude than that of Al. As a result, the electrons in Ga experience greater force of attraction by the nucleus than in Al and hence atomic radius of Ga is slightly less than that of Al. The ionic radii, however, follow a regular trend. 3. The first ionization enthalpies (IE1) of the elements of group 13 are lower than the corresponding elements of group 2, i.e., alkaline earth metals. This is due to the reason that elements of group 13 have three electrons in the valence shell – two of these are present in the s-orbital and one in the p-orbital. For the first ionization enthalpy (IE1), the electron has to be removed from the p-orbital in case of group 13 elements whereas in alkaline earth metals (group 2 elements), the s-electron of the same principal shell has to be removed. Since an s-electron is nearer the nucleus (more penetrating towards the nucleus), it is more strongly attracted than the p-electron of the same principal shell. Hence the removal of the p-electron is much easier than the s-electron. 4. The IE1 of Ga is only slightly higher (1 kJ mol-1) than that of Al while that of Tl is much higher than those of Al, Ga and In. This is due to the reason that Al follows immediately after s-block elements while Ga and In follow after d-block elements and Tl after d- and f-block elements. These extra d- and f-electrons do not shield (or screen) the outer shell-electrons from the nucleus very effectively. As a result, the valence electrons remain more tightly held by the nucleus and hence larger amount of energy is needed for their removal. This explains why Ga has higher ionization energy than Al. Further on moving down the group from Ga to In, the increased shielding effect (due to the presence of additional 4d-electrons) outweighs the effect of increased nuclear charge and hence the IE1 of In is lower than that of Ga. Thereafter, the effect of increased nuclear charge outweighs the shielding effect due the presence of additional 4f and 5d electrons and hence the IE1 of Tl is higher than that of In. 5. Boron and Aluminium which show an oxidation state of +3 only but Gallium, Indium and Thallium show oxidation states of both +1 and +3. Further, as we move down the group, the stability of the +3 oxidation state decreases while that of +1 oxidation state increases. For example, +1 oxidation state of Tl is more stable than +3. This is because as we move down the group, the tendency of s-electrons of the valence shell to participate in bond formation decreases. This reluctance of the s-electrons to participate in bond formation is called Inert pair effect. In other words, the ns2 electrons in Ga, In and Tl tends to remain paired. This is due to poor or ineffective shielding of the ns2 electrons of the valence shell by intervening d- and f- electrons. The inert pair effect becomes more predominant as we go down the group because of increased nuclear charge which outweighs the effect of the corresponding increase in atomic size. The s-electrons thus become more tightly held (more penetrating) and, therefore, become more reluctant to participate in bond formation. Thus down the group, +1 oxidation state becomes more and more stable as compared to +3 oxidation state. 6. The elements of group 13 are less electropositive or metallic as compared to alkali metals (group 1) and alkaline earth metals (group 2). On moving down the group, the electropositive character of the elements first increases from B to Al and then decreases from Al to Tl. Amongst the elements of group 13, B has the highest sum of first three ionization enthalpies, i.e., IE1 + lE2 + IE3. As a result, it has little tendency to lose electrons and hence is least electropositive amongst group 13 elements. In other words, as expected, it is a non-metal and a poor conductor of electricity. However, as we move from B to Al, the sum of IE1 + lE2 + IE3 decreases substantially (6857 kJ mol-1 to 5137 kJ mol-1) due to increase in the atomic size and hence Al has a high tendency to lose electrons. In other words, Al is highly electropositive. The electropositive character of the remaining elements can be more easily explained on the basis of their respective electrode potentials. Since the electrode potentials for the reaction, M3+ (aq) + 3e- = M (s) increases from Al to Tl, therefore, their electropositive character decreases accordingly. 7. The elements of boron family (group 13) are more electronegative than the elements of alkali metals (group 1) and alkaline earth metals (group 2). The electronegativity first decreases from B to Al and then increases marginally. This is because of the increase in the atomic size and consequently decreases in the attraction of nucleus towards the electrons. 8. The melting points of group 13 elements do not show a regular trend as shown by elements of groups 1 and 2. This is probably due to the unusual crystal structures of B and Ga. Actually, the melting points decrease sharply on moving down the group up to Ga and then increase from In to Tl. In fact, Ga is a liquid with an incredible low melting point of 303K. Boron has a high melting point (2453K) because its crystal structure consists of icosahedra units. In contrast, the crystal structure of Ga is quite different from that of B. The unique structure suggests that Ga consists of almost discrete diatomic molecules and hence its melting point is exceptionally low. Another unusual property of Ga is that like Ge and Bi, liquid Ga expands when it changes into solid, i.e. density of solid Ga is less than that of liquid Ga. In contrast, the elements Al, In and Tl have close packed structures. Their melting points decrease from Al to In and increases again for Tl. However, the boiling points of these elements follow a regular trend and decrease regularly on moving down the group. 9. Due to smaller atomic and ionic radii, the elements of group 13 have higher densities as compared to elements of group 2. On moving down the group, the densities increase. This is due to corresponding increase in the atomic mass of the element which outweighs the effect of increased atomic size. The densities of B and Al are, however, quite lower than those of other members. 10. Boron, because of its small size and high sum of first three ionization energies, does not lose its three valence electrons to form B3+ ions. Therefore, it does not form ionic compounds. Instead, boron forms mainly covalent compounds by sharing its valence electrons. On moving down the group, from B to Tl, the atomic size increases and ionization energies decrease, and the tendency of these elements to form covalent compounds decreases while that of ionic compounds increases. Unlike B3+, Al3+ ions have small size and high charge. Therefore, it has only a little tendency to form ionic compounds, but instead has a great tendency to form covalent compounds. Electronic configuration of boron in the excited state is 1s2 2s1 2px1 2py1. It undergoes sp2 hybridisation forming three half-filled hybrid orbitals which can form three covalent bonds. The 2pz orbital is empty which can accept a lone pair of electrons. Hence boron compounds are electron-deficient and act as Lewis acids. 11. The reducing power of the elements of group 13 decreases down the group. The electrode potential is a measure of the reducing power. Lower is the value of electrode potential, stronger is the reducing agent. On moving down the group, the value of electrode potential increases, so the reducing power decreases. 12. Since the elements of group 13 have similar valence shell electronic configurations, they exhibit similar physical and chemical properties. However, the difference in electronic configurations of heavier elements (Ga, In and Tl) from lighter elements (B and Al) influences the physical and chemical behavior. B and Al have ns2 np1 type configuration whereas the heavier elements have filled d- and f-orbitals in between the noble gas core and valence electrons. Because of this, the variation in properties amongst the members of group 13 elements are much more striking as compared to those in alkali metals and alkaline earth metals.

Some important chemical properties of boron family are described below: 1. Hydrides. The elements of group 13 (boron family) do not combine directly with hydrogen. However, a number of hydrides of the elements of this group have been prepared by indirect methods. Boron forms a number of covalent hydrides called boranes. These are of two types : nido-boranes having the general formula BnHn+4 and arachno-boranes having the formula BnHn+6. Out of all these boranes, diborane (B2H6) is the most important. Other members of group 13 also form a few stable hydrides but they are polymeric in nature, e.g. (AlH3)n, (GaH3)n and (InH3)n and contain M....H....M bridges (M = Al, Ga or In). However, their stability decreases as we move down the group from Al to Tl due to a corresponding decrease in the strength of the M —H bond as the size of the atom increases.

2. Halides. All the elements of group 13 form trihalides of the general formula MX3 where X = F, Cl, Br or I. Tl (III) iodide is however unknown. Due to small size and high ionization enthalpy, boron forms covalent trihalides. BF3 is gaseous, BCl3 and BBr3 are liquids, while BI3 is a solid. All these trihalides are planar molecules in which B is sp2 hybridized. Due to the presence of only six electrons in respective valence shells, all these trihalides are Lewis acids. The relative acid strength of the boron trihalides follows the sequence: BF3 < BCl3 < BBr3 < BI3. The order of the strength can be explained on the basis of the tendency of the halogen atom to back donate its lone pair of electrons to the boron atom through pπ - pπ bonding. Since the sizes of the vacant 2pz orbital of B and any of the 2p orbital of F containing a lone pair are almost identical, the lone pair of electrons on F is donated easily to the B atom. As the size of the halogen atom increases from Cl to Br to I, the extent of overlap between 2p orbital of B and a bigger np orbitaI of the halogen (3p in Cl, 4p in Br and 5p in I) decreases. It is interesting to note that both boron and aluminium halides are Lewis acids but only aluminium halides exist as dimers whereas boron halides exist only as monomers. This is because boron atom is so small that it cannot accommodate four large sized halogen atoms around it.

3. Coordination complexes. Group 13 elements form complexes much more rapidly than s-block elements because of their smaller size and higher charge. However, due to the absence of d-orbitals, B forms only tetrahedral complexes such as [BF4]−, [BH4]− etc, but due to the presence of d-orbitals, Al, Ga, In and Tl form octahedral complexes, such as [AlF6]3−, [GaCl6]3−, [InCl6]3−, and [TlCl6]3−. Similarly, all these elements form octahedral aqua ions, i.e. [M(H2O)6]3+ where M = Al, Ga, In and Tl. Incidentally, Al also forms tetrahedral complexes such as Li+[AlH4]−. Further, because of the small size and increased nuclear charge, these ions exert sufficient attraction on water molecules. Therefore, salts such as chlorides, sulphates, nitrates and perchlorates of Al, Ga, In and Tl exist as hydrates. Similarly, aluminium sulphate reacts with some alkali metal and NH4+ ions to form double salts of the formula, M2SO4.Al2(SO4)3.24H2O or MAl(SO4)2.12H2O where M = Na+, K+, Rb+ and NH4+. These double salts are commonly known as alums and are extensively used in the softening of hard water and as mordant (which helps to bind the dye to the fabric) in dyeing and printing of textiles.

Explain why? 1. Why Aluminium, though an electropositive metal, finds extensive use as a structural material. * Being electropositive, Al readily reacts with air to form a hard protective layer of alumina (AI2O3) which protects it from further action. It is because of this reason, that Al is extensively used as a structural material.

2. Gold has much higher first ionization energy than boron, yet gold is a metal while boron is a non-metal. Explain. * This is explained on the basis of their crystal structure. Gold has a co-ordination number of 12 while boron has a coordination number of only 4. 3. Why B−X bond distance in BX3 is shorter than theoretically expected value? * This is due to pπ - pπ back bonding of the fully filled p-orbital of halogen (X) into the empty p-orbital of boron.

Properties of Group 14 elements 1. The atomic radii of group 14 elements are smaller than the corresponding elements of group 13. This is because when we move from group 13 to group 14 within the same period, the effective nuclear charge increases and hence the atomic radius decreases. The atomic radii of group 14 elements regularly increase as we move down the group. It is due to addition of a new energy shell in each succeeding element. The increase in atomic radii from Si onwards is, however, small due to ineffective shielding of the valence electrons by the intervening d-and f- orbitals. 2. The first ionization enthalpies of group 14 elements are higher than those of the Corresponding group 13 elements. This is because of greater nuclear charge and smaller size of the atoms of group 14 elements. The ionization enthalpy decreases steadily on moving from carbon down the group to lead. The decrease is very sharp from carbon to silicon while there is a slight increase in the first ionization enthalpy value of Pb as compared to that of Sn. This slight increase in the value of ionization enthalpy from Sn to Pb is due to the combined effect of poor shielding of d-electrons in Ge and Sn and d- and f-electrons in Pb and the increased atomic size of the elements. 3. Carbon can take up four electrons to form carbide ion, C4− ion. However, such a process is energetically not favorable since the chemical species is highly charged and thus requires large amount of energy for adding four electrons. However, carbon forms some carbides such as Be2C, CaC2, SiC and Al4C3 in which carbon is supposed to be present either as C22− or C4− ion. 4. Like carbon, silicon also shows an oxidation state of +4. The remaining elements of this group, i.e. Ge, Sn and Pb, however, show two oxidation states of +2 and +4 due to inert pair effect which arises due to ineffective shielding of the valence s-electrons by the intervening d- and/or f-electrons. Evidently, as the number of d and or f-electrons increases down the group from Ge to Pb, the inert pair effect becomes more and more prominent. As a result, the stability of the + 4 oxidation state decreases while that of the +2 oxidation state increases from Ge to Pb. 5. Compounds of group 14 elements which show an oxidation state of +4 are expected to be covalent because of their extremely small size whereas compounds which show an oxidation state of +2 are expected to be ionic because of large size and small charge. For example, SnCl2 is ionic solid while SnCl4 is a covalent liquid. Further as we move down the group, the tendency of the elements to form covalent compounds decreases whereas the tendency to form ionic compounds increases. 6. Electropositive character—Metallic character. The group 14 elements are less electropositive and hence less metallic than the group 13 elements because of smaller atomic size and higher ionization energy from C to Pb. Thus, carbon is strictly non-metallic, silicon is essentially a non-metal, germanium is a semi-metal (metalloid) with pronounced metallic character while tin and lead are typical metals. 7. The elements of group 14 are more electronegative than group 13 elements because of smaller size. Electronegativity, however, decreases down the group because of increase in atomic size. 8. The melting and boiling points of group 14 elements (carbon family) are much higher than those of the corresponding elements of the group 13 elements (boron family). This is due to the reason that atoms of the elements of group 14 can form four bonds with each other and hence there exist strong binding forces between their atoms both in the solid as well as in the liquid states. As a result, their melting points and boiling points are higher than corresponding elements of group 13. The melting points and boiling points decrease as we move down the group due to a corresponding decrease in the inter-atomic forces of attraction. However, the melting point of tin is lower than that of lead. 9. Carbon has the remarkable property of catenation which may be defined as “the ability of like atoms to link with one another through covalent bonds”. This is due to smaller size and higher electronegativity of carbon atom. The property of catenation depends upon the strength of element—element bond. Since the bond energy C—C bond is very large (355 kJ mol−1), carbon forms long straight or branched C—C chains or rings of different sizes and shapes. However, as we move down the group, the bond energies decrease rapidly, viz., Si—Si (297 kJ mol−1), Ge—Ge (260 kJ mol−1), Sn—Sn (240 kJ mol−1) and Pb—Pb (81 kJ mol−1) and therefore the tendency of catenation decreases alongwith. 10. All the elements of this group, except lead, show allotropy. Carbon exists in a number of allotropic forms out of which the two main crystalline forms are diamond and graphite. In diamond, carbon is sp3 hybridized. Each carbon thus forms covalent bonds with four other carbon atoms which lie at the corners of a regular tetrahedron. As a result, diamond exists as a three dimensional network solid. On the other hand, in graphite each carbon atom is sp2 hybridized and is linked to three other carbon atoms by three single covalent bonds forming hexagonal layers. The additional p-electron from each carbon forms an extended delocalized π-bonding system encompassing the entire layer due to the ability of carbon to form pπ - pπ bonds among its atoms. Various layers are held together by weak van der Waals' forces. 11. Graphite is thermodynamically more stable than diamond since its free energy of formation is 1.9 kJ mol−1 at room temperature and atmospheric pressure. Although the conversion of diamond into graphite is thermodynamically favorable, yet it normally does not occur because of high energy of activation for the process. The reverse process, i.e., conversion of graphite into diamond is thermodynamically not possible but can be done only under forcing conditions. Thus, graphite can be converted into diamond at 1873 K under a pressure of 50,000—60,000 atmospheres. 12. Diamond has the highest thermal conductivity of any known substance (about five times that of Cu) although it is a bad conductor of electricity. It is because of its high thermal conductivity, diamond tipped tools do not over heat and hence are extensively used for drilling and cutting purposes. 13. “Bucky ball” or “Buckminster fullerene” is an allotrope of carbon which is obtained when graphite is vaporized with the help of high power laser. It is a spherical molecule having the molecular formula C60 and has 60 vertices, each vertex being occupied by a carbon atom. Like graphite, here also each carbon atom is sp2 hybridized. It has been named so because of its resemblance to a soccer ball. 14. Oil dag and Aqua dag. The suspension of graphite in oil which is used as a lubricant in heavy machinery is called oil dag. Similarly a colloidal dispersion of graphite in water is called aqua dag. 15. Due to small size and high electronegativity, carbon has a strong tendency is form pπ - pπ multiple bonds either with itself (C = C, C C) or with other atoms of similar size such as oxygen (C = O) and nitrogen (C = N, C N). However, as we move down the group from carbon to lead, this ability to form multiple bonds decreases drastically due to a corresponding increase in size and decrease in electronegativity of the atom. The reluctance of silicon to form such pπ - pπ bonds to it is shown by the following facts: elemental silicon exists only in the diamond structure and not in the graphite structure and no form of elemental silicon is comparable to graphite. Again, CO2 containing two C = O double bonds is a gas while SiO2 is a solid, which is due to an infinite three- dimensional network of Si—O single bonds. 16. Silicon and other heavier elements of group 14 can form dπ - pπ multiple bonds because of the presence of vacant d-orbitals in them. This tendency is particularly strong in case of silicon linked to oxygen and nitrogen. For example, the geometry around the nitrogen atom in trimethylamine, N(CH3)3, is pyramidal (N is sp3 hybridized), whereas in the case of similar silicon compound, N(SiH3)3, called trisilylamine, it is planar triangular (N is sp2 hybridized). The reason being that the lone pair of electrons in the 2p-orbital of N overlaps with the empty d-orbital of Si to form dπ - pπ bond. As a result, transference of electrons occurs from N to Si and hence N(SiH3)3 is a weaker base than N(CH3)3. 17. Carbon because of the absence of d-orbitals, cannot expand its valence shell and hence its maximum covalency or coordination number is four. However, Si, Ge, Sn and Pb, due to the availability of vacant d-orbitals, show a coordination of greater than 4 (i.e. 5 and 6) forming penta-coordinated and hexa-coordinated complexes. For example, [SiF5]-, [SiF6J2-, [GeCl6]2-, [Sn(OH)6]2-, [Pb(OH)6]2-, and [PbCl6]2- etc.

Some important chemical properties of carbon family are described below: 1. Hydrides. All the elements of group 14 form covalent hydrides of the type MH4 directly or indirectly. The number of hydrides, their stability and the ease of formation decreases down the group. For example, carbon forms a large number of cyclic and acyclic compounds with hydrogen called hydrocarbons. Silicon and germanium form comparatively very small number of hydrides of the general formula MnH2n+2 (M = Si, n = 1 to 8 and M = Ge, n = 1 to 5) called silanes and germanes respectively. Tin and lead form one hydride each, i.e., SnH4 (stannane) and PbH4 (plumbane). The stability of the hydrides of group 14 elements decreases and hence their reducing powder increases as we move from CH4 to PbH4 due to a corresponding decrease in the strength of the M—H bond as the size of the element M increases down the group from C to Pb. 2. Halides. All the elements of group 14 form tetrahalides of the general formula MX4. All these tetrahalides are essentially covalent compounds and have tetrahedral shapes. The thermal stability and ionic character of these halides decreases with the increasing atomic number or the size of the halogen atom. Thus PbCl4 and PbBr4 are unstable while PbI4 is unknown. 3. The tetrachloride of carbon (CCl4) is not hydrolyzed by water although the tetrachlorides of all other elements of this group are easily hydrolyzed. This is because carbon has no d-orbitals and hence cannot expand its coordination number beyond 4. However, silicon can expand its octet (coordination number beyond four) due to the availability of energetically suitable vacant d-orbitals in its atom where the attack of the lone pair of oxygen atom of water molecule, forming a coordinate bond between the central atom and oxygen, takes place. Afterwards, loss of HCl occurs and one Cl atom in SiCl4 is replaced by a —OH group. This process continues till all the four Cl atoms are replaced by —OH groups yielding Si(OH)4, i.e. silicic acid. 4. The tetrahalides of these elements except carbon also combine with halogen acids to form complex ions, e.g. SiF4 + 2HF → H2SiF6. 5. All the elements, except carbon and silicon, i.e. Ge, Sn and Pb also form dihalides, MX2. The stability of these dihalides increases steadily as we move down the group from Ge to Pb, i.e., GeX2