Words: 5311
Due to small size and high nomination enthalpy, the electrons of lithium cannot be emitted. 7. Alkali metals do not form deposition ions because of their very high second nomination enthalpies. 8. All alkali metals are paramagnetic but their salts are diamagnetic. 9.
Alkali metals impart characteristic lord to the flame, I. E. , lithium imparts crimson red, sodium imparts golden yellow, potassium imparts pale violet while rubidium and cesium 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 nomination 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 De 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. Lie does not form peroxides because the strong positive field around the lithium ion (Lie+) attracts the negative charge so strongly that it does not permit monoxide (0-) ion to combine with oxygen to form peroxide ion (022-). 11. The peroxide anions are diamagnetic and are oxidation agents. 12.
Nanny reacts with CA of the air and has men used in submarines and confined places, as it removes both CA and produces 02. Manama + CHIC -+ enhance +02 13. Supersedes contain the ion [02-] and are paramagnetic and colored due to the presence of an unpaired electron. Lila and Anna are yellow, KEY is orange. Orbs is brown and CSCW is orange. Supersedes are stronger oxidation agents than peroxides and give H2O and 02 on treatment with water or acids. Potassium superposed Inorganic Chemistry is used in breathing masks, submarines and space capsules because it removes CA and at the same time produces 02. KEY + CHIC -+ CHECKS + 302 KICK + + OH -+ CHECK + 302 14. Lithium cannot be stored 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. Lie is the exception. It is used as scavenger in the metallurgical operations to remove oxygen and nitrogen. It is the lightest meta now. Its least tussle, least dense and least soft of all alkali metals. It has highest specific heat among all the alkali metals and specific heat decreases from Lie to CSS. 5. Due to large negative electrode potentials, alkali metals are strong reducing agents. The reducing character increases from An to CSS. However, lithium is the strongest reducing agent among all the alkali metals in spite of its highest Nomination enthalpy (E). This is because of extensive hydration of Lie+ ions and large amount of energy released during hydration, more than compensates the higher II value of lithium. Alkali metals are better reducing agents than H2O. Therefore, they react with compounds containing acidic hydrogen atoms liberating H2O 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 : LICK KICK > Orbs > CLC. Low melting point of LICK as compared to NCAA is most probably due to the covalent nature of Lick. Covalent character of lithium halides follows the order: Ill > Libra > LICK > Life. 17. Solutions of alkali metals in liquid ammonia are highly conducting and deep blue in color.
This is because of nominated cationic 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 impound with nitrogen, I. E. Lien. Properties of Group 2 elements [Alkaline earth metals] 1 . Elements of group 2 (or ‘IA) are called alkaline earth, since: (I) Their oxides are alkaline in nature like alkali metal oxides, and, (it) the oxides of Ca, Sir and Baa are found in earth’s surface. The general electron configuration of alkaline earth metals is ins. The last alkaline earth metal, I. . , 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 table than M+ ions in spite of the fact that IEEE of alkaline earth metals is almost double than the III . 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 MGM) also impart characteristic colors to the flame, e. G. , Ca imparts brick red, strontium gives crimson and barium gives green color. Be and MGM do not impart any characteristic color to the flame because of their high ‘E. . 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 nomination 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: Mages > Cases > Users > Bases. 1. 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. 2. Beech, 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. . , 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 AY (143 pm).
This is due to the filling of electrons in d-orbital. In-between AY (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-orbital. As the d-orbital 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 AY. As a result, the electrons in Ga experience greater force of attraction by the nucleus than in AY and hence atomic radius of Ga is slightly less than that of AY.
The ionic radii, however, follow a regular trend. 3. The first nomination enthalpies (III) 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 nomination enthalpy (III), 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 III of Ga is only slightly higher (1 k mol-l) than that of AY while that of TTL is much higher than those of AY, Ga and In. This is due to the reason that AY follows immediately after s-block elements while Ga and In follow after d-block elements and TTL 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 nomination energy than AY. Further on moving down the group from Ga to In, the increased shielding effect (due to the presence of additional d-electrons) outweighs the effect of increased nuclear charge and hence the III of In is lower than that to Ga Thereafter, the detect to increased nuclear charge outweighs the shielding effect due the presence of additional if and ad electrons and hence the III of TTL is higher than that of In. . Boron and Aluminum 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 TTL 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 ins electrons in Ga, In and IT tends to remain paired. This is due to poor or ineffective shielding of the ins 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. . The elements of group 13 are less electrophoresis or metallic as compared to alkali metals (group 1) and alkaline earth metals (group 2). On moving down the group, the electrophoresis character of the elements first increases from B to AY and then decreases from AY to TTL. Amongst the elements of group 13, B has the highest sum of first three nomination enthalpies, I. . , III + IEEE + IEEE. As a result, it has little tendency to lose electrons and hence is least electrophoresis 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 AY, the sum of III + IEEE + IEEE decreases substantially (6857 k mol-l to 5137 k mol-l) due to increase in the atomic size and hence AY has a high tendency to lose electrons. In other words, AY is highly electrophoresis. The electrophoresis 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, MM+ (as) + e- = M (s) increases from AY to TTL, therefore, their electrophoresis character decreases accordingly. . 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 electronegative first decreases from B to AY 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 TTL. In fact, Ga is a liquid with an incredible low melting point of KICK. Boron has a high melting point (KICK) 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. Insist of solid Ga is less than that of liquid Ga. In contrast, the elements AY, n an TTL nave close packed structures. Their melting point TTS decrease trot AY to In and increases again for TTL. 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 AY are, however, quite lower than those of other members. 10. Boron, because of its small size and high sum of first three nomination energies, does not lose its three valence electrons to form 83+ 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 TTL, the atomic size increases and nomination energies decrease, and he tendency of these elements to form covalent compounds decreases while that of ionic compounds increases. Unlike BE+, AAA+ 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 ass ass pix pappy. It undergoes SSP habitations forming three half- filled hybrid orbital which can form three covalent bonds. The ups orbital is empty which can accept a lone pair of electrons. Hence boron compounds are electron- deficient and act as Lewis acids. 1 1 . 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 TTL) from lighter elements (B and AY) influences the physical and chemical behavior. B and AY have ins nip type configuration whereas the heavier elements have filled d- and f-orbital 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 borates. These are of two types : Indo-borates having the general formula Bonn+4 and archon-borates having the formula Bonn+6.
Out of all these borates, debonair (BAH) is the most important. Other members of group 13 also form a few stable hydrides but they are polymeric in nature, e. G. (Allah)n, (Gash)n and (LUHN)n and contain M…. H…. M bridges (M = AY, Ga or In). However, their stability decreases as we move down the group from AY to TTL due to a corresponding decrease in the strength of the M -??H bond as the size of the atom increases. . Halides. All the elements of group 13 form trailside of the general formula MIX where X F r or T ( ) iodide is however unknown.
Due to small size and high nomination enthalpy, boron forms covalent trailside. BUFF is gaseous, BC and Barb are liquids, while 813 is a solid. All these trailside are planar molecules in which B is SSP hybridism. Due to the presence of only six electrons in respective valence shells, all these trailside are Lewis acids. The relative acid strength of the boron trailside follows the sequence: BUFF Since the sizes of the vacant ups orbital of B and any of the up 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 CLC to BRB to l, the extent of overlap between up orbital of B and a bigger NP orbital of the halogen (up in CLC, up in BRB and up in l) decreases. It is interesting to note that both boron and aluminum halides are Lewis acids but only aluminum halides exist as dimmers 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-orbital, B forms only tetrahedral complexes such as [BUFF]-, [BAH]- etc, but due to the presence of d-orbital, AY, Ga, In and TTL form octahedral complexes, such as [BIFF]3-, [Cacao]3-, nonce]3-, and [ATTIC]3-.
Similarly, all these elements form octahedral aqua ions, I. E. [M(H2O)6]3+ where M = AY, Ga, In and T’. Incidentally, AY also forms tetrahedral complexes such as Lie+[Allah]-. Further, because of the small size and increased nuclear charge, these ions exert sufficient attraction on water molecules. Therefore, salts such as chlorides, sulfates, nitrates and percolates of AY, Ga, In and TTL exist as hydrates. Similarly, aluminum sulfate reacts with some alkali metal and NH+ ions to form double salts of the formula, or where M = An+, K+, Orb+ and NH+.
These double salts are commonly known as alums and are extensively used in the feting 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 Aluminum, though an electrophoresis metal, finds extensive use as a structural material. * Being electrophoresis, AY readily reacts with air to form a hard protective layer of alumina (AY) which protects it from further action. It is because of this reason, that AY is extensively used as a structural material. 2.
Gold has much higher first nomination energy than boron, yet gold is a metal while boron is a non-metal. Explain. * This is explained on the basis of their crystal Truckee. Gold has a co-ordination number of 12 while boron has a coordination number of only 4. 3. Why B-X bond distance in BOX is shorter than theoretically expected value? * This is due to pen – pen back bonding of the fully filled p-orbital of halogen (X) into the empty p-orbital of boron. Properties to 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 Is onwards is, however, small due to ineffective shielding of the valence electrons by the intervening d-and f- orbital. 2. The first nomination 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 nomination 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 nomination enthalpy value of BP as compared to that of Sin. This slight increase in the value of nomination enthalpy from Sin to BP is due to the combined effect of poor shielding of d-electrons in Ge and Sin and d- and f-electrons in BP and the increased atomic size of the elements. 3. Carbon can take up four electrons to form carbide ion, CA- 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 Beech, Ca, SIC and ACACIA in which carbon is supposed to be present either as CA- or CA- ion. 4. Like carbon, silicon also shows an oxidation state of +4. The remaining elements of this group, I. E. Ge, Sin and BP, 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 BP, 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 BP. 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, Syncs is ionic solid while Syncs 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. . Electrophoresis character-??Metallic character. The group 14 elements are less electrophoresis and hence less metallic than the group 13 elements because of smaller atomic size and higher nomination energy from C to BP. 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. Electronegative, however, decreases down the group because of increase in atomic size. 8. The letting 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 Witt each other and hence there exist strong binding tortes 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 electronegative 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 k mol-l), 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, biz. Is-??Is (297 k mol-l), Ge-??Ge (260 k mol-l), sin-??sin (240 k mol-l) and BP-??BP (81 k mol-l) and therefore the tendency of catenation decreases longhair. 10. All the elements of this group, except lead, show allotrope. Carbon exists in a number of allotropic forms out of which the two main crystalline forms are diamond and graphite.
In diamond, carbon is SSP hybridism. 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 ACH carbon atom is SSP hybridism 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 idealized II-bonding system encompassing the entire layer due to the ability of carbon to form pen – pen bonds among its atoms.
Various layers are held together by weak van deer Walls’ forces. 11. Graphite is thermodynamically more stable than diamond since its free energy of formation is 1. 9 k mol-l at room temperature and atmospheric pressure. Although the conversion of diamond into graphite is thermodynamically favorable, yet it normally goes 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 atmospheres. 2. Diamond has the highest thermal conductivity of any known substance (about five times that of Cue) 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. Buck ball” or “Businessmen fulfiller” 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 CA and has 60 vertices, each vertex being occupied by a carbon atom.
Like graphite, here also each carbon atom is SSP hybridism. It has been named so because of its resemblance to a soccer ball. 14. Oil tag and Aqua tag. The suspension of graphite in oil which is used as a lubricant in heavy machinery is called oil tag. Similarly a colloidal dispersion of graphite in water is called aqua tag. 15. Due to small size ND high electronegative, carbon has a strong tendency is form pen – pen 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, TN TTY to tort multiple bonds decreases drastically due to a corresponding increase in size and decrease in electronegative of the atom. The reluctance of silicon to form such pen – pen 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, CA containing two C = O double bonds is a gas while Isis is a solid, which is due to an infinite three- dimensional network of Is-??O single bonds. 16.
Silicon and other heavier elements of group 14 can form den – pen multiple bonds because of the presence of vacant d-orbital 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 thermodynamic, N(CHI)3, is pyramidal (N is SSP hybridism), whereas in the case of similar silicon compound, N(Shih)3, called transatlantic, it is planar triangular (N is SSP hybridism). The reason being that the lone pair of electrons in the up- orbital of N overlaps with the empty d-orbital of Is to form den – pen bond.
As a result, transference of electrons occurs from N to Is and hence N(Shih)3 is a weaker base than N(CHI)3. 17. Carbon because of the absence of d-orbital, cannot expand its valence shell and hence its maximum covalently or coordination number is four. However, S’, Ge, Sin and BP, due to the availability of vacant d-orbital, show a coordination of greater than 4 (I. E. 5 and 6) forming pentane-coordinated and hexane- coordinated complexes. For example, [Sift]-, [Skiffs-, [Gecko]2-, ND [Pubic]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 MPH 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 Mann+2 (M = Is, n = 1 to 8 and M =Ge, n = 1 to 5) called silages and Germans respectively. Tin and lead form one hydride each, I. . , Shin (stagnate) and EBPP (plumbing).
The stability of the hydrides of group 14 elements decreases and hence their reducing powder increases as we move from CHI to EBPP 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 BP. 2. Halides. All the elements of group 14 form terrestrials of the general formula MIX. All these terrestrials 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 f the halogen atom.
Thus Pubic and BRB are unstable while EBPP is unknown. 3. The tetrachloride of carbon (CHIC) is not hydrolysis by water although the tetrachloride of all other elements of this group are easily hydrolysis. This is because carbon has no d-orbital 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-orbital in its atom where the attack of the lone pair of oxygen atom of water molecule, forming a coordinate bond twine the central atom and oxygen, takes place.
Afterwards, loss of HCI occurs and one CLC atom in Sick is replaced by a -??OH group. This process continues till all TN four CLC atoms are replaced by -??OH groups yielding Is(OH)4, I. E. Silica acid. 4. The terrestrials of these elements except carbon also combine with halogen acids to form complex ions, e. G. Sift + UHF -??+ Hashish. 5. All the elements, except carbon and silicon, I. E. Ge, Sin and BP also form idealizes, MIX. The stability of these idealizes increases steadily as we move down the group from Ge to BP, I. E. , Gees