chemistry notes for neet chapter 18.pdf

60 Chapter E3 18 s and p-Block Elements The group 1 of the periodic table contains six elements, namely lithium (Li), sodium (Na), potassium (K), rubidium (Rb), caesium (Cs) and francium (Fr). All these elements are typical metals. Francium is radioactive U with longest lived isotope 223 Fr with hal...

60 Chapter E3 18 s and p-Block Elements The group 1 of the periodic table contains six elements, namely lithium (Li), sodium (Na), potassium (K), rubidium (Rb), caesium (Cs) and francium (Fr). All these elements are typical metals. Francium is radioactive U with longest lived isotope 223 Fr with half life period of only 21 minute. These are usually referred to as alkali metals since their hydroxides form strong bases or alkalies. D YG (1) Electronic configuration Elements Discovery 3 Li 11 Na 19 K 37 Rb 87 Fr Electronic configuration ( ns 1 ) Arfwedson (1817) [He]2 2 s 1 Davy (1807) [Ne]10 3 s 1 Davy (1807) [Ar]18 4 s 1 Bunsen (1861) [Kr]36 5 s 1 Bunsen (1860) [Xe] 54 6 s 1 U 55 Cs Percy (1939) [Rn]86 7 s 1 ST (2) Occurrence : Alkali metals are very reactive and thus found in combined state some important ores of alkali metals are given ahead. (i) Lithium : Triphylite, Petalite, lepidolite, Spodumene [LiAl(SiO ) ], Amblygonite [Li(Al F)PO ] (ii) Sodium : Chile salt petre (NaNO ), Sodium chloride (NaCl), Sodium sulphate (Na SO ), Borax (Na B O 10H O), Glauber salt (Na SO.10H O) (iii) Potassium : Sylime (KCl), carnallite (KCl.MgCl.6H O) and Felspar (K O.Al O.6SiO ) (iv) Rubidium : Lithium ores Lepidolite, triphylite contains 0.7 to 3% 3 3 4 2 2 2 3 4 (ii) Being highly electropositive in nature, it is not possible to apply the method of displacing them from their salt solutions by any other element. (iii) The aqueous solutions of their salts cannot be used for extraction by electrolytic method because hydrogen ion is discharged at cathode instead of an alkali metal ions as the discharge potentials of alkali metals are high. However, by using Hg as cathode, alkali metal can be deposited. The alkali metal readily combines with Hg to form an amalgam from which its recovery difficult. The only successful method, therefore, is the electrolysis of their fused salts, usually chlorides. Generally, another metal chloride is added to lower their fusion temperature. ID Alkali Metals and Their Compounds 3 2 4 7 2 2 2 4 2 2 2 Rb O 2 (v) Caesium : Lepidolite, Pollucite contains 0.2 to 7% Cs O (3) Extraction of alkali metals : Alkali metals cannot be extracted by the usual methods for the extraction of metals due to following reasons. (i) Alkali metals are strong reducing agents, hence cannot be extracted by reduction of their oxides or other compounds. 2 fusion Fused NaCl : NaCl    Na   Cl – Electrolysis : Anode : 2Cl  Cl 2  2e  of fused salt: Cathode : 2 Na   2e  2 Na (4) Alloys Formation (i) The alkali metals form alloys among themselves as well as with other metals. (ii) Alkali metals also get dissolved in mercury to form amalgam with evolution of heat and the amalgamation is highly exothermic. Physical properties (1) Physical state (i) All are silvery white, soft and light solids. These can be cut with the help of knife. When freshly cut, they have bright lustre which quickly tarnishes due to surface oxidation. (ii) These form diamagnetic colourless ions since these ions do not have unpaired electrons, (i.e. M has ns configuration). That is why alkali metal salts are colourless and diamagnetic. (2) Atomic and ionic radii (i) The alkali metals have largest atomic and ionic radii than their successive elements of other groups belonging to same period. (ii) The atomic and ionic radii of alkali metals, however, increases down the group due to progressive addition of new energy shells. No doubt the nuclear charge also increases on moving down the group but the influence of addition of energy shell predominates + Li Atomic radius (pm) Ionic radius of M+ 152 60 0 Na K Rb 186 95 227 133 248 148 Cs Fr 265 375 169 – Li Na K Rb Cs Fr 2 7 2 7 4 1- 4 (7) Hydration of Ions (i) Hydration represents for the dissolution of a substance in water to get adsorb water molecule by weak valency force. Hydration of ions is the exothermic process (i.e energy is released during hydration) when ions on dissolution water get hydration. (ii) The energy released when 1 mole of an ion in the gaseous state is dissolved in water to get it hydrated is called hydration energy M (g)   Aq M  (aq) ; H  – ve. (iii) Smaller the cation, greater is the degree of hydration. Hydration energy is in the order of, Li > Na > K > Rb > Cs (iv) Li being smallest in size has maximum degree of hydration and that is why lithium salts are mostly hydrated, LiCl. 2H O also lithium ion being heavily hydrated, moves very slowly under the influence of electric field and, therefore, is the poorest conductor current among alkali metals ions It may, therefore, be concluded that it is the degree of hydration as well as the size of ion is responsible for the current carried by an ion. + + + + + + 2 Cs+ > Rb+ > K+ > Na+ > Li+ Li+ > Na+ > K+ > Rb+ > Cs+ Relative conducting power Cs+ > Rb+ > K+ > Na + > Li+ Relative ionic radii Relative hydrated ionic radii (8) Electronegativity, Electro positivity and metallic character. (i) These metals are highly electropositive and thereby possess low values of electronegativities. Metallic character and electro positivity increase from Li to Cs (Li < Na < K < Rb < Cs) (ii) Electronegativity of alkali metals decreases down the group as the trend of numerical values of electronegativity given below on Pauling scale suggests. IE1 IE2 Na K Rb 418 3069 403 2650 Cs Fr D YG Li Ionisation energy U ID melting point (K) 453.5 370.8 336.2 312.0 301.5 – boiling point (K) 1620 1154.4 1038.5 961.0 978.0 – (ii) The lattice energy of these atoms in metallic crystal lattice relatively low due to larger atomic size and thus possess low melting point and boiling point on moving down the group, the atomic size increases and binding energy of their atoms in crystal lattice decreases which results lowering of melting point. (iii) Lattice energy decreases from Li to Cs and thus melting point and boiling also decreases from Li to Cs. (5) Ionisation energy and electropositive or metallic character (i) Due to unpaired lone electron in ns sub-shell as well as due to their larger size, the outermost electron is far from the nucleus, the removal of electron is easier and these low values of ionisation energy. (I.E.) (ii) Ionisation energy of these metal decreases from Li to Cs. 2- 2 60 0 K Cr O is orange because of orange coloured Cr O ion, KMnO is violet because of violet coloured MnO ion. E3 ions (pm) (3) Density (i) All are light metals, Li, Na and K have density less than water. Low values of density are because these metals have high atomic volume due to larger atomic size. On moving down the group the atomic size as well as atomic mass both increase but increase in atomic mass predominates over increase in atomic size or atomic volume and therefore the ratio mass/volume i.e. density gradually increases down the groups (ii) The density increases gradually from Li to Cs, Li is lightest known metal among all. Li = 0.534, Na = 0.972, K = 0.86, Rb = 1.53 and Cs = 1.87 g/ml at 20 C. (iii) K is lighter than Na because of its unusually large atomic size. (iv) In solid state, they have body centred cubic lattice. (4) Melting point and Boiling point (i) All these elements possess low melting point and boiling point in comparison to other group members. 520 495 7296 4563 376 2420 – – Li Electronegativity 0.98 Na 0.93 Re moval of Re movalof Li :1s 2 2 s1  Li  :1s 2  Li 2  : 1s1 2 s electron 1 s electron Removal of 1s electrons from Li and that too from completely filled configuration requires much more energy and a jump in 2nd ionisation is noticed. (iii) Lower are ionisation energy values, greater is the tendency to lose ns electron to change in M ion (i.e. M M +e ) and therefore stronger is electropositive character. (iv) Electropositive character increases from Li to Cs. Due to their strong electropositive character, they emit electrons even when exposed to light showing photoelectric effect. This property is responsible for the use of Cs and K in photoelectric cell. + + – ST 1 U + (6) Oxidation number and valency (i) Alkali metals are univalent in nature due to low ionisation energy values and form ionic compounds. Lithium salts are, however, covalent. (ii) Further, the M ion acquires the stable noble gas configuration. It requires very high values of energy to pull out another electron from next to outer shell of M ion and that is why their second ionisation energy is very high. Consequently, under ordinary conditions, it is not possible for these metals to form M ion and thus they show +1 oxidation state. (iii) Since the electronic configuration of M ions do not have unpaired electron and thus alkali metal salts are diamagnetic and colourless. Only those alkali metal salts are coloured which have coloured anions e.g. + + 2+ Rb Cs 0.82 Fr 0.79 – Fr being radioactive elements and thus studies on physical properties of this element are limited. (9) Specific heat : It decreases from Li to Cs. Li A jump in 2nd ionisation energy (huge difference) can be explained as, K 0.82 Specific heat (Cal/g) Na 0.941 0.293 0.17 K Rb Cs 0.08 0.049 – Fr (10) Conduction power : All are good conductors of heat and electricity, because of loosely held valence electrons. (11) Standard oxidation potential and reduction properties (i) Since alkali metals easily lose ns electron and thus they have high values of oxidation potential i.e., 1 M  aq M  (aq)  e (ii) The standard oxidation potentials of a alkali metals (in volts) are listed below, Li +3.05 Na K +2.71 +2.93 Rb +2.99 Cs +2.99 (iii) More is oxidation potential, more is the tendency to get oxidized and thus more powerful is reducing nature in aqueous medium that is why alkali metals liberate H from H O and HCl. 2 2 2 H 2O  2 M 2 MOH  H 2 ; 2 HCl  2 M 2 MCl  H 2 (iv) However, an examination of ionisation energy for alkali metals reveals that Li should have the minimum tendency to lose electron and thus its reducing nature should be minimum. The greatest reducing nature of Li in aq. medium is accounted due to the maximum hydration energy of Li ion. For Lithium + Li(s) Li(g) ; Li(g) Li + Li  (g)  (g) Li H = Heat of sublimation, H 1  e ; H = IE  (aq); 2 s 1 H = – Heat of hydration, H 3 h whereas Na does so vigorously, K reacts producing a flame and Rb, Cs do so explosively. Li(s)  H 2O Li (aq )  e ; H  H1  H 2  H 3  H s  IE1  H h Similarly, for sodium, h + : Li+ < Na+ < K+ < Rb+ < Cs+ (1) Formation of oxides and hydroxides (i) These are most reactive metals and have strong affinity for O quickly tranish in air due to the formation of a film of their oxides on the surface. These are, therefore, kept under kerosene or paraffin oil to protect them from air, M  O2 M 2 O  M 2 O2 2 K  O2 D YG 1 O 2 Li 2 O ; 2 Na  O 2 Na 2 O 2 2 Lithuim oxide 2 Li  KO 2 Potassium super oxide The reactivity of alkali metals towards oxygen to form different oxides is due to strong positive field around each alkali metal cation. Li being smallest, possesses strong positive field and thus combines with small anion O to form stable Li O compound. The Na and K being relatively larger thus exert less strong positive field around them and thus reacts with + 2– + + 2 larger oxygen anion i.e, O22 and O12  to form stable oxides. The monoxide, peroxides and superoxides have O and O 22  , O 12 ions respectively. The structures of each are,............. O:] : O...... O : [ x O.. O x]2 [:O U 2 1– – Monoxide (O ) 2 Peroxide (O ) 2– 2 Superoxide (O ) – 2 The O ion has a three electron covalent bond and has one electron unpaired. It is therefore superoxides are paramagnetic and coloured KO is light yellow and paramagnetic substance. (iii) The oxides of alkali metals and metal itself give strongly alkaline solution in water with evolution of heat –1 ST 2 1 H2; 2 Li2O  H 2O 2 LiOH; M  H 2 O MOH  2 H  ve H  ve Na 2 O 2  2 H 2 O 2 NaOH  H 2 O 2(l) ; 2 KO 2  2 H 2 O 2 KOH  H 2 O2(l)  O2(g) ; H  ve The peroxides and superoxides act as strong oxidising agents due to formation of H O (iv) The reactivity of alkali metals towards air and water increases from Li to Cs that is why lithium decomposes H O very slowly at 25 C 2 1 H2  e 2 (vi) Alkali metals also form hydrides like NaBH , LiAlH which are good reducing agent. At anode: H  – 4 2 o 2 4 (3) Carbonates and Bicarbonates (i) The carbonates (M CO ) & bicarbonates (MHCO ) are highly stable to heat, where M stands for alkali metals. (ii) The stability of these salts increases with the increasing electropositive character from Li to Cs. It is therefore Li CO decompose on heating, Li CO Li O+CO (iii) Bicarbonates are decomposed at relatively low temperature, 2 3 3 2 2 3 2 3 2 0 300 C 2 MHCO 3    M 2 CO 3  H 2 O  CO 2 (iv) Both carbonates and bicarbonates are soluble in water to give alkaline solution due to hydrolysis of carbonate ions or bicarbonate ions. (4) Halides (i) Alkali metals combine directly with halogens to form ionic halide MX. (ii) The ease with which the alkali metals form halides increases from Li to Cs due to increasing electropositive character from Li to Cs. (iii) Lithium halides however have more covalent nature. Smaller is the cation, more is deformation of anion and thus more is covalent nature in compound. Also among lithium halides, lithium iodide has maximum covalent nature because of larger anion which is easily deformed by a cation. Thus covalent character in lithium halides is, LiI > LiBr > LiCl > LiF (iv) These are readily soluble in water. However, lithium fluoride is sparingly soluble. The low solubility of LiF is due to higher forces of attractions among smaller Li and smaller F ions (high lattice energy). (v) Halides having ionic nature have high m.pt. and good conductor of current. The melting points of halides shows the order, NaF > NaCl > NaBr > Nal (vi) Halides of potassium, rubidium and caesium have a property of combining with extra halogen atoms forming polyhalides. + H  ve 2  U 2 2 2 – NaH fused Contains Na and H i.e., + 2 2  At cathode: Na +e Na; Peroxide (ii) When burnt air (O ), lithium forms lithium oxide (Li O) sodium forms sodium peroxide (Na O ) and other alkali metals form super oxide (Mo i.e. KO ,RbO or CsO ) 2 2 ID 2 2 2 2 2 Chemical properties Oxide 2M+ H 2MH ; Reactivity towards H is Cs < Rb < K < Na < Li. (iii) The metal hydrides react with water to give MOH and H ; MH + H O MOH + H (iv) The ionic nature of hydrides increases from Li to Cs because of the fact that hydrogen is present in the these hydrides as H and the smaller cation will produce more polarisation of anion (according to Fajans rule) and will develop more covalent character. (v) The electrolysis of fused hydrides give H at anode. 2  released : Li+ > Na+ > K+ > Rb+ > Cs+ Frequency released : Li+ < Na+ < K+ < Rb+ < Cs+ E3 Energy released – 60 H for Li > H for Na. Therefore, large negative H values are observed in case of Li and this explains for more possibility of Li to get itself oxidized or have reducing nature. (12) Characteristic flame colours : The alkali metals and their salts give characteristic colour to Bunsen flame. The flame energy causes and excitation of the outermost electron which on reverting back to its initial position gives out the absorbed energy as visible light. These colour differ from each other Li –crimson, Na–Golden yellow, K – Pale violet , Rb-Red violet and Cs –Blue violet. These different colours are due to different ionisation energy of alkali metals. The energy released is minimum in the case of Li and increases in the order. h 1 H2 2 (v) The basic character of oxides and hydroxides of alkali metals increases from Li to Cs. This is due to the increase in ionic character of alkali metal hydroxides down the group which leads to complete dissociation and leads to increase in concentration of OH ions. (2) Hydrides (i) These metals combine with H to give white crystalline ionic hydrides of the general of the formula MH. (ii) The tendency to form their hydrides, basic character and stability decreases from Li to Cs since the electropositive character decreases from Cs to Li. M  H 2 O MOH  Na(s)  H 2 O Na  (aq)  e ; H  H(s)  IE1  H h – KI + I KI ; In KI the ions K and I are present + 2 3 3(aq) – 3 3 3 + 3 3 Na  (x  y)NH 3 [ Na(NH 3 )x ]  [e(NH 3 )y ] Ammoniated cation Ammoniated electron (iv) It is the ammoniated electron which is responsible for blue colour, paramagnetic nature and reducing power of alkali metals in ammonia solution. However, the increased conductance nature of these metals in ammonia is due to presence of ammoniated cation and ammonia solvated electron. (v) The stability of metal-ammonia solution decreases from Li to Cs. (vi) The blue solution on standing or on heating slowly liberates hydrogen, 2M + 2NH 2MNH + H. Sodamide (NaNH ) is a waxy solid, used in preparation of number of sodium compounds. (6) Nitrates : Nitrates of alkali metals (MNO ) are soluble in water and decompose on heating. LiNO decomposes to give NO and O and rest all give nitrites and oxygen. 3 2 2 2 3 2 2 2MNO 2MNO + O (except Li) 3 2 2 4 LiNO 2Li O + 4NO + O 3 2 2 2 4 2 4 D YG 2 2 4 2 4 0 ; 2K + H 2KH 300 C   2NaH 2Na + H  2 2Na + Cl 2NaCl ; 2K + Cl 2KCl 2 2 ; 2K + S K S U 2Na + S Na S 2 2 3Na + P Na P ; 3K + P K P (ii) Li reacts, however directly with carbon and nitrogen to form carbides and nitrides. 2Li + 2C LiC ; 6Li + 2N 2 Li N (iii) The nitrides of these metals on reaction with water give NH. M N + 3H O 3MOH + NH (9) Reaction with acidic hydrogen : Alkali metals react with acids and other compounds containing acidic hydrogen (i.e, H atom attached on F,O, N and triply bonded carbon atom, for example, HF, H O, ROH, RNH , CH  CH) to liberate H. ST 3 2 3 2 3 2 2 3 3 3 3 2 2 M  H 2 O MOH  1 1 H 2 ; M  HX MX  H 2 2 2 M  ROH ROH  1 1 H 2 ; M  RNH 2 RNHNa  H 2 2 2 2 2 2 2 2  2 LiOH  Li 2 O  H 2 O (9) LiHCO is liquid while other metal bicarbonates are solid. (10) Only Li CO decomposes on heating 3 3 (8) Reaction with non-metals (i) These have high affinity for non-metals. Except carbon and nitrogen, they directly react with hydrogen, halogens, sulphur, phosphorus etc. to form corresponding compounds on heating. 2 2 U 2 4 2 2 3 2 (7) Sulphates (i) Alkali metals’ sulphate have the formula M SO. (ii) Except Li SO , rest all are soluble in water. (iii) These sulphates on fusing with carbon form sulphides, M SO + 4C M S + 4CO (iv) The sulphates of alkali metals (except Li) form double salts with the sulphate of the trivalent metals like Fe, Al, Cr etc. The double sulphates crystallize with large number of water molecules as alum. e.g. KSO. Al (SO ). 24 HO. 2 2 ID 3 (10) Complex ion formation : A metal shows complex formation only when it possesses the following characteristics, (i) Small size (ii) High nuclear charge (iii) Presence of empty orbitals in order to accept electron pair ligand. Only Lithium in alkali metals due to small size forms a few complex ions Rest all alkali metals do not possess the tendency to form complex ion. Anomalous behaviour of Lithium Anomalous behaviour of lithium is due to extremely small size of lithium its cation on account of small size and high nuclear charge, lithium exerts the greatest polarizing effect out of all alkali metals on negative ion. Consequently lithium ion possesses remarkable tendency towards solvation and develops covalent character in its compounds. Li differs from other alkali metals in the following respects, (1) It is comparatively harder than other alkali metals. Li can’nt be stored in kerosene as it floats to the surface, due to its very low density. Li is generally kept wrapped in parrafin wax. (2) It can be melted in dry air without losing its brilliance. (3) Unlike other alkali metals, lithium is least reactive among all. It can be noticed by the following properties, (i) It is not affected by air. (ii) It decomposes water very slowly to liberate H. (iii) It hardly reacts with bromine while other alkali metals react violently. (4) Lithium is the only alkali metal which directly reacts with N to form Lithium nitride (Li N) (5) Lithium when heated in NH forms amide, Li NH while other metals form amides, MNH. (6) When burnt in air, lithium form Li O sodium form Na O and Na O other alkali metals form monoxide, peroxide and superoxide. (7) Li O is less basic and less soluble in water than other alkali metals. (8) LiOH is weaker base than NaOH or KOH and decomposes on heating. 60 3 E3 (5) Solubility in liquid NH (i) These metals dissolve in liquid NH to produce blue coloured solution, which conducts electricity to an appreciable degree. (ii) With increasing concentration of ammonia, blue colour starts changing to that of metallic copper after which dissolution of alkali metals in NH ceases. (iii) The metal atom is converted into ammoniated metal in i.e. M (NH ) and the electron set free combines with NH molecule to produce ammonia solvated electron. 2 3 heat Li 2 CO 3   Li 2 O  CO 2. Na CO , K CO etc. do not decompose on heating. (11) LiNO and other alkali metal nitrates give different products on 2 3 2 3 3 heating 4LiNO = 2Li O+4NO + O ; 2NaNO = 2NaNO + O (12) LiCl and LiNO are soluble in alcohol and other organic solvents. These salts of other alkali metals are, however, insoluble in organic solvents. (13) LiCl is deliquescent while NaCl, KBr etc. are not. Lithium chloride crystals contain two molecules of water of crystallisation ( LiCl. 2H O). Crystals of NaCl KBr, KI etc do not conation water of crystallisation. (14) Li SO does not form alums like other alkali metals. (15) Li reacts with water slowly at room temperature Na reacts vigorously Reaction with K. Rb and Cs is violent. (16) Li reacts with Br slowly. Reaction of other alkali metals with Br is fast. (17) Li CO Li C O , LiF , Li PO are the only alkali metal salts which are insoluble or sparingly soluble in water. Diagonal Relationship of Li with Mg Due to its small size lithium differs from other alkali metals but resembles with Mg as its size is closer to Mg Its resemblance with Mg is known as diagonal relationship. Generally the periodic properties show either increasing or decreasing trend along the group and vice versa along 3 2 2 2 3 2 2 3 2 2 4 2 2 3 2 2 4 2 3 4 2 Li CO Li O + CO ; Mg CO MgO + CO (5) Hydroxides and nitrates of both Li and Mg decompose on heating to give oxide. Hydroxides of both Li and Mg are weak alkali. 3 2 3 2 2 2 2 2 2 2 2 2 2KNO 2KNO + O (6) Both Li and Mg combine directly with N to give nitrides Li N and Mg N. Other alkali metals combine at high temperature, 6Li + N 2Li N; 3Mg + N Mg N. Both the nitrides are decomposed by water to give 2 2 2 3 2 3 NH 2 3 2 3 Li N + 3H O 3LiOH + NH ; 2 3 D YG 3 Mg N + 6H O 3Mg(OH) + 2NH (7) Bicarbonates of Li and Mg are more soluble in water than 3 2 2 carbonates whereas carbonates of alkali metals are more soluble. (8) Both Li and Mg combine with carbon on heating. 2Li + 2C Li C ; Mg + 2C Mg C (9) The periodic properties of Li and Mg are quite comparable 2 2 Li Mg ST U Electronegativity 1.0 1.2 Atomic radii 1.23 1.36 Ionic radii 0.60(Li ) 0.65(Mg ) Atomic volume 12.97 c.c 13.97 c.c (10) Both have high polarizing power. Polarizing Power = Ionic charge / (ionic radius). (11) Li and Mg Form only monooxide on heating in oxygen. + +2 2 4Li + O 2 Li O ; 2Mg + O 2 MgO (12) Li SO like MgSO does not form alums. (13) The bicarbonates of Li and Mg do not exist in solid state, they exist in solution only. (14) Alkyls of Li and Mg (R. Li and R.MgX) are soluble in organic solvent. (15) Lithium chloride and MgCl both are deliquescent and separate out from their aqueous solutions as hydrated crystals, LiCl. 2H O and MgCl. 2H O. Uses of Lithium (1) It is used as a deoxidiser in metallurgy of Cu and Ni. 2 2 4 2 At anode : Cl  Cl  e  ; Cl  Cl Cl 2  0 NaCl because  (–0.83V) is more than E Na / Na (–2.71V). Anode and cathode are separated by means of a wire gauze to prevent the reaction between Na and Cl 2. (3) Compound of sodium (i) Sodium chloride : It is generally obtained by evaporation of sea water by sun light. However NaCl so obtained contains impurities like CaSO 4 , CaCl 2 and MgCl2 which makes the salt deliquescent. It is then purified by allowing HCl gas to pass through the impure saturated solution of NaCl. The concentration of Cl  ions increases and as a result pure NaCl gets precipitated due to common ion effect. (ii) Sodium hydroxide NaOH (Caustic soda) Preparation (a) Gossage process : Na 2CO 3 (10 % solution)  Ca(OH )2 2 NaOH  CaCO 3 (b) Electrolytic method : Caustic soda is manufactured by the electrolysis of a concentrated solution of NaCl. At anode: Cl  discharged; At cathode: Na  discharged (c) Castner - Kellener cell (Mercury cathode process) : NaOH obtained by electrolysis of aq. solution of brine. The cell comprises of rectangular iron tank divided into three compartments. Outer compartment – Brine solution is electrolysed ; Central compartment – 2% NaOH solution and H 2 Properties : White crystalline solid, highly soluble in water, It is only sparingly soluble in alcohol. (a) Reaction with salt : 2 4 2 2 2 At cathode : Na   e  Na ; 3 2 2 NaCl ⇌ Na   Cl . U 2 3 (2) Extraction of sodium : It is manufactured by the electrolysis of fused sodium chloride in the presence of CaCl 2 and KF using graphite anode and iron cathode. This process is called Down process. ID 2 3 sodium borate, (Na 2 B4 O7. 10 H 2 O). EH0 2 O / H 2 2LiOH Li O + H O ; Mg(OH) MgO + H O Hydroxides of other alkali metals are stable towards heat while their nitrates give O and nitrite. 2 petre), Na 2 SO 4.10 H 2 O (Glauber's salt), borax (sodium tetraborate or Sodium cannot be extracted from aqueous 2 2Mg(NO ) 2MgO + 4NO + O 3 (1) Ores of sodium : NaCl (common salt), NaNO 3 (chile salt 2 4 LiNO 2Li O + 4NO + O 3 Sodium and its compounds E3 2 (2) It is used as an alloying metal with (i) Pb to give toughened bearings. (ii) Al to give high strength Al-alloy for aircraft industry. (iii) Mg (14% Li) to give extremely tough and corrosion resistant alloy which is used for armour plate in aerospace components. 60 the period which brought the diagonally situated elements to closer values. Following are the characteristic to be noted. Period Group I Group II 2 Li Be Na Mg 3 (1) Both Li and Mg are harder and higher m.pt than the other metals of their groups. (2) Due to covalent nature, chlorides of both Li and Mg are deliquescent and soluble in alcohol and pyridine while chlorides of other alkali metals are not so. (3) Fluorides, phosphates of Li and Mg are sparingly soluble in water whereas those of other alkali metals are soluble in water. (4) Carbonates of Li and Mg decompose on heating and liberate CO Carbonates of other alkali metals are stable towards heat and decomposed only on fusion. FeCl3  3 NaOH Fe(OH )3 (Insoluble hydroxide)   3 NaCl HgCl2  2 NaOH 2 NaCl  Hg(OH )2 H 2 O  HgO  yellow unstable AgNO3  2 NaOH 2 NaNO 3  2 AgOH Ag2 O   H 2 O Brown 2 Zn, Al, Sb , Pb, Sn and As forms insoluble hydroxide which dissolve in excess of NaOH (amphoteric hydroxide). Na 2 CO 3  H 2 O H 2 CO 3  2 Na   2OH  heat NH 4 Cl  NaOH   NaCl  NH 3   H 2O Weak acid (b) Reaction with halogens : (c) It is readily decomposed by acids with the evolution of CO 2 X 2  2 NaOH (cold) NaX  NaXO  H 2 O gas. sod. hypohalite 5 NaX  NaXO3  3 H 2O ; 3 X 2  6 NaOH (hot) (d) Na 2 CO 3  H 2 O  CO 2 2NaHCO 3 (Sod. halate) Uses : In textile and petroleum refining, Manufacturing of glass, NaOH soap powders etc. (iv) Sodium peroxide (Na O ) Preparation : It is manufactured by heating sodium metal on aluminium trays in air (free from CO 2 ) (X  Cl, Br, I) (c) Reaction with metals : Weakly electropositive metals like Zn, Al 2 and Sn etc. 60 Zn  2 NaOH Na 2 ZnO2  H 2   Na O  2 Na  O 2 (air)  2 2 (d) Reaction with sand, SiO : 2 Na 2 SiO3  Sod. silicate(glass) Properties : (a) When pure it is colourless. The faint yellow colour of commercial product is due to presence of small amount of superoxide (NaO 2 ). H 2O (e) Reaction with CO: (b) On coming with moist air it become white due to formation of NaOH and Na 2 CO 3. o 150  200 C NaOH  CO   HCOONa 5 10 atm E3 2 NaOH  SiO2 2 Sod. formate 2 Na 2 O 2  2 H 2 O 4 NaOH  O 2 ; 2 NaOH  CO 2 Na 2 CO 3  H 2 O (c) It is powerful oxidising agent. It oxidises Cr (III) hydroxide to sodium chromate, Mn (II) to sodium manganate and sulphides to sulphates. Uses : As a bleaching agent and it is used for the purification of air in confined spaces such as submarines because it can combine with CO 2 D YG (iii) Sodium carbonate or washing soda, Na 2 CO 3 U ID NaOH breaks down the proteins of the skin flesh to a pasty mass, therefore it is commonly known as caustic soda. Caustic property : sodium hydroxide breaks down the proteins of the skin flesh to a pasty mass, therefore, it is commonly known as caustic soda. Uses : Sodium hydroxide is used : (a) in the manufacture of soidum metal, soap (from oils and fats), rayon, paper, dyes and drugs, (b) for mercurinzing cotton to make cloth unshrinkable and (c) as a reagent in the laboratory. It exists in various forms, namely anhydrous sodium carbonate Na CO (soda-ash); monohydrate Na 2 CO 3.H 2 O (crystal carbonate); hyptahydrate 2 2 Na 2 CO 3.7 H 2 O and decahydrate Na 2 CO 3.10 H 2 O (washing soda or sal soda). Preparation : (a) Solvay process : In this process, brine (NaCl) , NH 3 and CO 2 are the raw materials. NH 3  CO 2  H 2 O NH 4 HCO 3 U oC 30 NH 4 HCO 3  NaCl   NaHCO 3   NH 4 Cl oC 250 2 NaHCO 3    Na 2 CO 3  H 2 O  CO 2 ST 2 NH 4 Cl  Ca(OH )2 CaCl 2  2 H 2 O  2 NH 3 slaked lime CaCl 2 so formed in the above reaction is a by product of solvay (v) Micro cosmic salt [Na (NH ) HPO. 4H O] 4 4 2 Prepared by dissolving equimolar amounts of Na 2 HPO4 and NH 4 Cl in water in 1 : 1 ratio followed by crystallization NH 4 Cl  Na 2 HPO4  Na( NH 4 )HPO4  NaCl  Crystallization Na(NH 4 )HPO4.4 H 2 O (Colourless crystal) Chemical properties : On heating M.C.S, NaPO3 is formed. NaPO3 forms coloured beads with oxides of transition metal cloudy SiO 2  Na( NH 4 )HPO4  NaPO3  H 2 O  NH 3 (Sodium meta phosphate) NaPO3  (Trans parent glassy bead)  CuO  CuNaPO 4 (blue bead) NaPO3  CoO  CoNaPO4 (blue bend) NaPO3  MnO  NaMnO4 (blue bead) process. Properties dry air (a) Na 2 CO 3.10 H 2 O   Na 2 CO 3.H 2 O  9 H 2 O (decahydra te) (Monohydrate)  Na 2CO 3. H 2O   to give Na 2 CO 3 and oxygen, 2CO 2  2 Na 2 O 2 2 Na 2 CO 3  O 2. Na 2CO 3 It does not decompose on funrther heating even to redness (m.pt. 853°C) (b) It is soluble in water with considerable evolution of heat. Uses : (a) For the formation of sodium meta phosphate and copper sodium phosphate (b) It is used for the detection of colured ion (c) It is espacially used for testing silica with which a cloudy bead containing floating properties of silica is obtained. (vi) Sodium bi Carbonate (NaHCO , Baking soda) Preparation : It is an inter mediate compound in manufacture of sodium carbonate by the solvay’s process 3 (v) Radium : Pitch blende (U O ); (Ra in traces); other radium rich minerals are carnotite [K UO )] (VO ) 8H O and antamite [Ca(UO ) ] (3) Extraction of alkaline earth metals (i) Be and Mg are obtained by reducing their oxides carbon, NaCl  NH 3  CO 2  H 2  NaHCO 3  NH 4 Cl 3 2 o 50 100 C Properties: 2 NaHCO 3    Na 2 CO 3  H 2 O  CO 2 It is amphiprotic HCO 3  H  ⇌ H 2 CO 3  ⇌H  2 8 2 2 2 BeO + C Be + CO ; MgO + C Mg + CO CO 32  Uses : (a) Baking powder contains NaHCO 3 , Ca(H 2 PO4 )2 and starch. Improved Baking powder contains 40% starch 30% NaHCO 3 , (ii) The extraction of alkaline earth metals can also be made by the reduction of their oxides by alkali metals or by electrolysing their fused salts. (4) Alloy formation : These dissolve in mercury and form amalgams. Physical properties (1) Physical state : All are greyish-white, light, malleable and ductile metals with metallic lustre. Their hardness progressively decrease with increase in atomic number. Although these are fairly soft but relatively harder than alkali metals. 60 20% NaAl(SO 4 )2 and 10% CaH 2 (PO4 ) (b) In pharmacentical industry (Antacids etc.) (c) Fire extingerishers. (vii) Sodium Sulphate Na SO or salt cake 2 4 4 Preparation : It is the by-product of HCl industry 2 NaCl  H 2 SO 4  Na 2 SO 4  HCl Properties : When aqueous solution of Na 2 SO 4 is cooled below o 32 C Glauber’s salt (Na 2 SO 4.10 H 2 O) gets crystallised and if cooled to 12 C , Na 2 SO 4 7 H 2 O crystals are formed. (indry air) Na 2 SO 4.10 H 2 O   Na 2 SO 4  10 H 2 O Uses : Na 2 SO 4 finds use in paper industry detergent and glass manufacturing. Be Mg Ca Sr Ba Ra 112 31 160 65 197 99 215 113 222 135 – 140 (ii) The atomic radii of alkaline earth metals are however smaller than their corresponding alkali metal of the same period. This is due to the fact that alkaline earth metals possess a higher nuclear charge than alkali metals which more effectively pulls the orbit electrons towards the nucleus causing a decrease in size. U Alkaline Earth Metals and Their Compounds Atomic radius (pm) Ionic radius of M2+ ion (pm) ID o (2) Atomic and ionic radii (i) The atomic and ionic radii of alkaline earth metals also increase down the group due to progressive addition of new energy shells like alkali metals. E3 HCO 3 2 D YG The group 2 of the periodic table consists of six metallic elements. These are beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr), barium (Ba) and radium (Ra). These (except Be) are known as alkaline earth metals as their oxides are alkaline and occur in earth crust. (1) Electronic configuration Element Electronic configurations ( ns 2 ) [He] 2 s 2 12 Mg [ Ne ] 3 s 2 20 Ca [ Ar] 4 s 2 38 Sr [Kr] 5 s 2 56 Ba 88 Ra U Be [Rn]7 s 2 Radium was discovered in the ore pitch blende by madam Curie. It is radioactive in nature. (2) Occurrence : These are found mainly in combined state such as oxides, carbonates and sulphates Mg and Ca are found in abundance in nature. Be is not very abundant, Sr and Ba are less abundant. Ra is rare element. Some important ores of alkaline earth metals are given below, (i) Baryllium : Beryl (3BeO.Al O.6SiO ); Phenacite (Be SiO ) (ii) Magnesium : Magnesite (MgCO ); Dolomite (CaCO. MgCO ); Epsomite(MgSO. 7H O); Carnallite (MgCl.KCl. 6H O); Asbestos [CaMg (SiO ) ] (iii) Calcium : Limestone (CaCO ); Gypsum : (CaSO.2H O), Anhydrite (CaSO ); Fluorapatite [(3Ca (PO ) CaF )] Phosphorite rock [Ca (PO ) ] (iv) Barium : Barytes (BaSO ) ; witherite (BaCO ) 2 3 2 2 4 3 4 3 3 2 3 2 3 2 4 3 4 3 4 2. 4 2 4 Be 1.84 Mg Ca Sr Ba Ra 1.74 1.55 2.54 3.75 6.00 (ii) The alkaline earth metals are more denser, heavier and harder than alkali metal. The higher density of alkaline earth metals is due to their smaller atomic size and strong intermetallic bonds which provide a more close packing in crystal lattice as compared to alkali metals. (4) Melting point and Boiling point (i) Melting points and boiling points of alkaline earth metals do not show any regular trend. Be [ Xe ] 6 s 2 ST 4 (3) Density (i) Density decreases slightly upto Ca after which it increases. The decrease in density from Be to Ca might be due to less packing of atoms in solid lattice of Mg and Ca. 2 3 3 4 2 Mg Ca Sr Ba Ra melting points (K) 1560 920 1112 1041 1000 973 boiling point (K) 2770 1378 1767 1654 1413 – (ii) The values are, however, more than alkali metals. This might due to close packing of atoms in crystal lattice in alkaline earth metals. (5) Ionisation energy and electropositive or metallic character (i) Since the atomic size decreases along the period and the nuclear charge increases and thus the electrons are more tightly held towards nucleus. It is therefore alkaline earth metals have higher ionisation energy in comparison to alkali metals but lower ionisation energies in comparison to p-block elements. (ii) The ionisation energy of alkaline earth metals decreases from Be to Ba. Be Mg Ca Sr Ba First ionisation energy (k J mol-1) 899 737 590 549 503 509 Second ionisation energy (kJ mol-1) 1757 1450 1146 1064 965 979 Ra (iii) The higher values of second ionisation energy is due to the fact that removal of one electron from the valence shell, the remaining electrons are more tightly held in which nucleus of cation and thus more energy is required to pull one more electron from monovalent cation. (iv) No doubt first ionisation energy of alkaline earth metals are higher than alkali metals but a closer look on 2nd ionisation energy of alkali metals and alkaline earth metals reveals that 2nd ionisation energy of alkali metals are more Be 520 1757 899 removal of 2 s removal of 1 s electron electron Electronegativity removal of 2 s electron electron The removal of 2 electron from alkali metals takes place from 1s sub shell which are more closer to nucleus and exert more nuclear charge to hold up 1s electron core, whereas removal of 2nd electron from alkaline earth metals takes from 2s sub shell. More closer are shells to the nucleus, more tightly are held electrons with nucleus and thus more energy is required to remove the electron. (v) All these possess strong electropositive character which increases from Be to Ba. (vi) These have less electropositive character than alkali metals as the later have low values of ionisation energy. (6) Oxidation number and valency (i) The IE of the these metals are much lower than IE and thus it appears that these metals should form univalent ion rather than divalent ions but in actual practice, all these give bivalent ions. This is due to the fact that M ion possesses a higher degree of hydration or M ions are extensively hydrated to form [M(H O) ] , a hydrated ion. This involves a large amount of energy evolution which counter balances the higher value of second ionisation energy. M M , H = IE + E M + H O [M(H O) ] ; H = – hydration energy. (ii) The tendency of these metals to exist as divalent cation can thus be accounted as, (a) Divalent cation of these metals possess noble gas or stable configuration. (b) The formation of divalent cation lattice leads to evolution of energy due to strong lattice structure of divalent cation which easily compensates for the higher values of second ionisation energy of these metals. (c) The higher heats of hydration of divalent cation which accounts for the existence of the divalent ions of these metals in solution state. (7) Hydration of ions (i) The hydration energies of alkaline earth metals divalent cation are much more than the hydration energy of monovalent cation. 1 2+ 2+ 2+ D YG x 2 2+ x ST U 2 Mg+ Hydration energy or Heat of hydration (kJ mol–1) 353 Mg2+ 1906 The abnormally higher values of heat of hydration for divalent cations of alkaline earth metals are responsible for their divalent nature. MgCl formation occurs with more amount of heat evolution and thus MgCl is more stable. (ii) The hydration energies of M ion decreases with increase in ionic radii. 2 2 2+ Be2+ Heat of hydration kJ mol–1 2382 Mg2+ 1906 Be Mg Ca Sr Ba 1.57 1.31 1.00 0.95 0.89 Be Mg 1.69 2.35 Ca2+ 1651 1484 Ca Sr Ba 2.87 2.89 2.90 (ii) All these metals possess tendency to lose two electrons to give M ion and are used as reducing agent. (iii) The reducing character increases from Be to Ba, however, these are less powerful reducing agent than alkali metals. (iv) Beryllium having relatively lower oxidation potential and thus does not liberate H from acids. 2+ 2 (11) Characteristic flame colours The characteristic flame colour shown are : Ca - brick red; Sr – crimson ; Ba-apple green and Ra- crimson. Chemical Properties (1) Formation of oxides and hydroxides (i) The elements (except Ba and Ra) when burnt in air give oxides of ionic nature M O which are crystalline in nature. Ba and Ra however give peroxide. The tendency to form higher oxides increases from Be to Ra. U 1 2 2+ ID nd x 2 (10) Standard oxidation potential and reducing properties (i) The standard oxidation potential (in volts) are, Li2+ : 1s1 removal of 2 s 2+ 2. (8) Electronegativities (i) The electronegativities of alkaline earth metals are also small but are higher than alkali metals. (ii) Electronegativity decreases from Be to Ba as shown below, Be : 1s2 , 2s2   Be+ : 1s2, 2s1   Be2+ : 1s2 1 2 (9) Conduction power : Good conductor of heat and electricity. Li : 1s2, 2s1   Li +: 1s2   2+ 2 2+ This may be explained as, 2 2 60 7296 Li 2 E3 1st ionisation energy (kJ mol–1) 2nd ionisation energy (kJ mol–1) of alkali metals e.g MgCl and CaCl exists as Mg Cl.6H O and CaCl 6H O which NaCl and KCl do not form such hydrates. (iv) The ionic mobility, therefore, increases from Be to Ba , as the size of hydrated ion decreases. Sr2+ Ba2+ 1275 (iii) Heat of hydration are larger than alkali metals ions and thus alkaline earth metals compounds are more extensively hydrated than those 2+ 2- 2M + O 2MO 2 (M is Be, Mg or Ca ) (M is Ba or Sr) 2M + O MO (ii) Their less reactivity than the alkali metals is evident by the fact that they are slowly oxidized on exposure to air, However the reactivity of these metals towards oxygen increases on moving down the group. (iii) The oxides of these metals are very stable due to high lattice energy. (iv) The oxides of the metal (except BeO and MgO) dissolve in water to form basic hydroxides and evolve a large amount of heat. BeO and MgO possess high lattice energy and thus insoluble in water. (v) BeO dissolves both in acid and alkalies to give salts i.e. BeO possesses amphoteric nature. 2 2 BeO+2NaOHNa BeO +H O ; BeO+2HClBeCl +H O 2 2 2 Sod. beryllate 2 2 Beryllium chloride (vi)The basic nature of oxides of alkaline earth metals increases from Be to Ra as the electropositive Character increases from Be to Ra. (vii)The tendency of these metal to react with water increases with increase in electropositive character i.e. Be to Ra. (viii) Reaction of Be with water is not certain, magnesium reacts only with hot water, while other metals react with cold water but slowly and less energetically than alkali metals. (ix) The inertness of Be and Mg towards water is due to the formation of protective, thin layer of hydroxide on the surface of the metals. (x) The basic nature of hydroxides increase from Be to Ra. It is because of increase in ionic radius down the group which results in a decrease in strength of M –O bond in M –(OH) to show more dissociation of hydroxides and greater basic character. (xi) The solubility of hydroxides of alkaline earth metals is relatively less than their corresponding alkali metal hydroxides Furthermore, the 2 solubility of hydroxides of alkaline earth metals increases from Be to Ba. Be (OH) and Mg (OH) are almost insoluble, Ca (OH) (often called lime water) is sparingly soluble whereas Sr(OH) and Ba (OH) (often called baryta water) are more soluble. The trend of the solubility of these hydroxides depends on the values of lattice energy and hydration energy of these hydroxides. The magnitude of hydration energy remains almost same whereas lattice energy decreases appreciably down the group leading to more –Ve values for H down the group. 2 2 2 2 2 solution H = H solution lattice energy + H (4) Halides (i) The alkaline earth metals combine directly with halogens at appropriate temperatures forming halides, MX. These halides can also be prepared by the action of halogen acids (HX) on metals, metal oxides, hydroxides and carbonates. M + 2HX MX + H ; MO + 2HX MX + H O M(OH) + 2HX MX +2H O MCO + 2HX MX + CO + H O Beryllium chloride is however, conveniently obtained from oxide 2 2 2 2 3 hydration energy More negative is H more is solubility of compounds. (xii) The basic character of oxides and hydroxides of alkaline earth metals is lesser than their corresponding alkali metal oxides and hydroxides. (xiii) Aqueous solution of lime water [Ca(OH) ] or baryta water [Ba(OH)] are used to qualitative identification and quantative estimation of carbon dioxide, as both of them gives white precipitate with CO due to formation of insoluble CaCO or BaCO 2 2 2 2 2 2 2 870 1070 K BeO  C  Cl 2   BeCl 2  CO solution 2 3 3 Ca(OH)2 + CO2 CaCO3 + H2O ; Ba(OH)2 + CO2 BaCO3 + H2O (white ppt) (white ppt) SO also give white ppt of CaSO and BaSO on passing through lime water or baryta water. However on passing CO in excess, the white 2 3 3 2 3 2 2 3 2 2 2 2 2 4 2 2 2 4 2 2 D YG 2 2+ 2 – 3 2 2 2+ 2 2 2 2 2 U 2 3 3 2 2 (aq) 4 2 (aq) 2 2 3 (aq) 3 3 ST 2 2 (g) 3(s) 2 (l) 2 3(s) Heat MCO3   MO  CO 2 (aq) 2 2 2 CaCl2 Brick red colour 2 SrCl2 BaCl2 Crimson colour Grassy green colour Structure of BeCl : In the solid phase polymeric chain structure with three centre two electron bonding with Be-Cl-Be bridged structure is shown below, 2 Cl Cl 202 PM 98o Be 82o Cl 263 pm Be Be Cl Cl – Cl In the vapour phase it tends to form a chloro-bridged dimer which dissociates into the linear triatomic monomer at high temperature at nearly 1200 K. (5) Solubility in liquid ammonia : Like alkali metals, alkaline earth metals also dissolve in liquid ammonia to form coloured solutions When such a solution is evaporated, hexammoniate, M(NH ) is formed. 3 6 (6) Nitrides (i) All the alkaline earth metals direct combine with N give nitrides, 3 4 2 2 2 2 2(s) 2 2 2 – – 2 increases. The evidence is provided by the following facts, (a) Beryllium chloride is relatively low melting and volatile whereas BaCl has high melting and stable. (b) Beryllium chloride is soluble in organic solvents. (iii) The halides of the members of this group are soluble in water and produce neutral solutions from which the hydrates such : MgCl 6H O, CaCl.6H O. BaCl 2H O can be crystallised. The tendency to form hydrated halides decreases with increasing size of the metal ions. (iv) BeCl is readily hydrolysed with water to form acid solution, BeCl + 2H O Be (OH) + 2HCl. (v) The fluorides are relatively less soluble than the chlorides due to high lattice energies. Except BeCl and MgCl the chlorides of alkaline earth metals impart characteristic colours to flame. U 2 2 ID turbidity of insoluble carbonates dissolve to give a clear solution again due to the formation of soluble bicarbonates, CaCO H O + CO Ca(HCO ) (2) Hydrides (i) Except Be, all alkaline earth metals form hydrides (MH ) on heating directly with H. M+ H MH. (ii) BeH is prepared by the action of LiAlH On BeCl 2BeCl + LiAlH 2BeH + LiCl + AlCl. (iii) BeH and MgH are covalent while other hydrides are ionic. (iv) The ionic hydrides of Ca, Sr, Ba liberate H at anode and metal at cathode. CaH Ca + 2H fusion Anode : 2H H + 2e Cathode : Ca + 2e Ca (v) The stability of hydrides decreases from Be to Ba. (vi) The hydrides having higher reactivity for water, dissolves readily and produce hydrogen gas. CaH + 2H O Ca(OH) + 2H  (3) Carbonates and Bicarbonates (i) All these metal carbonates (MCO ) are insoluble in neutral medium but soluble in acid medium. These are precipitated by the addition of alkali metal or ammonium carbonate solution to the solution of these metals. (NH ) CO + CaCl 2NH Cl + CaCO Na CO + BaCl 2NaCl + BaCO (ii) Alkaline earth metal carbonates are obtained as white precipitates when calculated amount of carbon dioxide is passed through the solution of the alkaline metal hydroxides. M(OH) + CO MCO + H O and sodium or ammonium carbonate is added to the solution of the alkaline earth metal salt such as CaCl. CaCl + Na CO CaCO +2 NaCl (iii) Solubility of carbonates of these metals also decreases downward in the group due to the decrease of hydration energy as the lattice energy remains almost unchanged as in case of sulphates. (vi) The carbonates of these metals decompose on heating to give the oxides, the temperature of decomposition increasing from Be to Ba. Beryllium carbonate is unstable. 60 2 2 2 E3 2 (ii) BeCl is essentially covalent, the chlorides MgCl , CaCl , SrCl and BaCl are ionic; the ionic character increases as the size of the metal ion 2 MN. 3 2 (ii) The ease of formation of nitrides however decreases from Be to Ba. (iii) These nitrides are hydrolysed water to liberate NH , M N + 6H O 3M(OH) + 2NH 3 3 2 2 2 3 (7) Sulphates (i) All these form sulphate of the type M SO by the action of H SO on metals, their oxides, carbonates or hydroxides. M + H SO MSO + H ; MO+H SO MSO +H O MCO + H SO MSO + H O+CO M(OH) + H SO MSO + 2H O (ii) The solubility of sulphates in water decreases on moving down the group BeSO and MgSO are fairly soluble in water while BaSO is completely insoluble. This is due to increases in lattice energy of sulphates down the group which predominates over hydration energy. (iii) Sulphate are quite stable to heat however reduced to sulphide on heating with carbon. 4 2 4 3 4 2 4 2 2 4 2 2 4 2 4 4 4 4 2 4 4 2 2 2 4 MSO + 2C MS+2CO 4 2 (8) Action with carbon : Alkaline metals (except Be, Mg) when heated with carbon form carbides of the type MC These carbides are also called acetylides as on hydrolysis they evolve acetylene. 2 MC + 2H OM(OH) + C H 2 2 2 2 (9) Action with sulphur and phosphorus : Alkaline earth metals directly combine with sulphur and phosphorus when heated to form sulphides of the type MS and phosphides of the type M P respectively. 3 2 M + S MS ; 3M + 2P M P Sulphides on hydrolysis liberate H S while phosphides on hydrolysis 3 2 2 evolve phosphine. 3 2 3 2 3 2MS + 2H O M(OH) + M(HS) 2 2 2 (10) Nitrates : Nitrates of these metals are soluble in water On heating they decompose into their corresponding oxides with evolution of a mixture of nitrogen dioxide and oxygen. 1 M ( NO 3 )2 MO  2 NO 2   O 2 2 – 4 Anomalous behaviour of Beryllium Beryllium differs from rest of the alkaline earth metals on account of its small atomic size, high electronegativity Be exerts high polarizing effect on anions and thus produces covalent nature in its compounds. Following are some noteworthy difference of Be from other alkaline earth metals, (1) Be is lightest alkaline earth metal. (2) Be possesses higher m.pt. and b.pt than other group members. (3) BeO is amphoteric in nature whereas oxides of other group members are strong base. (4) It is not easily effected by dry air and does not decompose water at ordinary temperature. (5) BeSO is soluble in water. D YG 2+ 4 (6) Be and Mg carbonates are not precipitated by (NH 4 )2 CO 3 in presence of NH Cl. U 4 (7) Be and Mg salts do not impart colour to flame. (8) Be does not form peroxide like other alkaline earth metals. ST (9) It does not evolve hydrogen so readily from acids as other alkaline earth metals do so. (10) It has strong tendency to form complex compounds. (11) Be N is volatile whereas nitrides of other alkaline earth metals are non-volatile. (12) It’s salts can never have more than four molecules of water of crystallization as it has only four available orbitals in its valence shell. (13) Berylium carbide reacts water to give methane whereas magnesium carbide and calcium carbide give propyne and acetylene respectively. 2 Be C+4H O2Be(OH) + CH Mg C + 4H O 2Mg(OH) + C H CaC + 2H O Ca(OH) + C H Diagonal relationship of Be with Al 2 2 3 4 2 2 3 2 3 2 2 2 2 3 4 6 2 3 2 2 and alluminates. Be + 2NaOH Na BeO +H 2Al + 6NaOH 2Na AlO + 3H (9) Be C and Al C both give CH on treating with water. Be C+ 2H OCH + 2BeO Al C + 6H O3CH + 2Al O (10) Both occur together in nature in beryl ore, 3BeO. Al O 6SiO. 2 4 2 2 4 2 4 2 3 2 3 3 4 ID 3 2 2 4 2 3 2 3. 2 (11) Unlike other alkaline earths but like aluminium, beryllium is not easily attacked by air (Also Mg is not attacked by air) (12) Both Be and Al react very slowly with dil. HCl to liberate H. (13) Both Be and Al form polymeric covalent hydrides while hydrides of other alkaline earth are ionic. (14) Both BeCl and AlCl are prepared is similar way. U 2- 4 Al O + 2NaOH 2NaAlO + H O Al O + 6HCl 2AlCl + 3H O (8) Be and Al both react with NaOH to liberate H forming beryllates 3 2+ 1- 3 2 (11) Formation of complexes (i) Tendency to show complex ion formation depends upon smaller size, high nuclear charge and vacant orbitals to accept electron. Since alkaline metals too do not possess these characteristics and thus are unable to form complex ion. (ii) However, Be on account of smaller size forms many complex such as (BeF ) , (BeF ). 2 3+ 60 2 2 2+ – Sulphides are phosphorescent and are decomposed by water 2 3+ 3 MS + dil. acid H S ; M P + dil. acid PH 3 2+ E3 2 Due to its small size Be differs from other earth alkaline earth metals but resembles in many of its properties with Al on account of diagonal relationship. (1) Be and Al have almost same and smaller size and thus favour for covalent bonding. (2) Both these form covalent compounds having low m. pt and soluble in organic solvent. (3) Both have same value of electronegativity (i.e. 1.5). (4) The standard O.P of these elements are quite close to each other ; Be =1.69 volts and Al = 1.70 volts. (5) Both become passive on treating with conc. HNO in cold. (6) Both form many stable complexes e.g. (BeF ) , (AlH ). (7) Like BeO, Al O is amphoteric in nature. Also both are high melting point solids. 2 2 3 BeO+ C+ Cl BeCl + CO 2 2 Al O + 3C +3Cl 2AlCl + 3CO (15) Both BeCl and AlCl are soluble in organic solvents and act as 2 3 2 3 2 3 catalyst in Friedel –Crafts reaction. (16) Both Be (OH) and Al (OH) are amphoteric whereas hydroxides of other alkaline earths are strong alkali. (17) The salts of Be and Al are extensively hydrated. (18) BeCl and AlCl both have a bridged polymeric structure. (19) Be and Al both form fluoro complex ions [BeF ] and [AlF ] in solution state whereas other members of 2nd group do not form such complexes. 2 2 3 3 2– 4 3– 6 Magnesium and its compounds (1) Ores of magnesium : Magnesite (MgCO3 ), Dolomite (MgCO3.CaCO 3 ) , Epsomite (epsom salt) (MgSO4.7 H 2 O) Carnallite (MgCl2. KCl. 6 H 2 O) Asbestos (CaMg3 (SiO3 )4 ), Talc (Mg2 (Si2 O5 )2. Mg(OH )2 ). (2) Extraction of magnesium : It is prepared by the electrolysis of fused magnesium chloride which in turn is obtained from carnallite and magnesite. Carnallite (MgCl2.KCl.6 H 2 O) can’t be directly converted into anhydrous MgCl2 by heating because all the water of crystallisation cannot be removed by heating. Moreover, strong heating may change it to MgO.  MgCl2  2 H 2 O  MgO  2 HCl  H 2 O In Dow’s process, magnesium chloride is obtained from sea water as MgCl2.6 H 2 O. It is rendered anhydrous by heating it in a current of dry of anhydrous MgCl2 ) and then electrolysed at 700 o C. (3) Compounds of magnesium (i) Magnesia (MgO) : It is used as magnesia cement. It is a mixture of MgO and MgCl2. It is also called Sorel's cement. (ii) Magnesium hydroxide : It aqueous suspension is used in Medicine as an antacid. Its medicinal name is milk of magnesia. (iii) Magnesium sulphate or Epsom salt (MgSO4. 7 H 2 O) : It is isomorphous with ZnSO 4. 7 H 2 O. It is used as a purgative in medicine, as a mordant in dyeing and as a stimulant to increase the secretion of bile. (iv) Magnesium chloride (MgCl2.6 H 2 O) : It is a deliquescent solid. Hydrated salt on heating in air undergoes partial hydrolysis. Heat MgCl2. 6 H 2 O    Mg(OH )Cl  HCl  5 H 2 O. oC 1 1 120 CaSO 4. 2 H 2 O    CaSO 4. H 2 O  1 H 2 O 2 Gypsum Plaster of 2 paris Plaster of paris : 1 Hardening H 2O H 2 O   CaSO 4. 2 H 2 O   CaSO 4.2 H 2 O 2 Setting orthorhomb ic Monoclinic(gypsum) Plaster of paris CaSO 4. o 200 C CaSO 4.2 H 2 O    CaSO 4 (anhydrous) Gypsum dead burnt plaster Gypsum when heated to about 200 o C is converted into anhydrous calcium sulphate. The anhydrous form (anhydrite) is known as dead burnt plaster because it does not set like plaster of paris when moistened with water. 60 HCl gas. The anhydrous magnesium chloride is fused with NaCl (to provide conductivity to the electrolyte and to lower the fusing temperature (v) Calcium Hydroxide Ca(OH )2 (slaked lime) CaO  H 2 O Ca(OH )2 (1) Ores of calcium : Lime stone or marble or chalk (CaCO 3 ), Gypsum (CaSO 4. 2 H 2 O), Dolomite (CaCO 3. MgCO3 ), Suspension of Ca(OH )2 in water is called milk of lime. Ca(OH)2  Cl 2 CaOCl 2  H 2 O (vi) Cement : (a) It is essentially a mixture of lime stone and clay. It is also called Portland cement because in presence of water it sets to a hard stone-like mass resembling with the famous Portland rock, a famous building stone of England. The approximate composition of cement is ID Fluorspar (CaF2 ), phosphorite Ca 3 (PO4 )2. Calcium phosphate is a constituent of bones and teeth. (2) Manufacture : It is manufactured by the electrolysis of a molten mixture of calcium chloride containing some calcium fluoride. Calcium chloride is obtained as a by product of the solvay process. (3) Compounds of calcium (i) Calcium oxide or Quick lime or Burnt lime (CaO) : It's aqueous suspension is known as slaked lime. E3 Ca(OH )2  CO 2 CaCO3  Ca(HCO3 )2 Calcium and its compounds Calcium oxide (CaO) 50 – 60 % 20 – 25% Alumina ( Al 2 O 3 ) 5 – 10% Magnesia (MgO) 1 – 3% Ferric oxide (Fe 2 O 3 ) 1 – 3% U Silica (SiO 2 ) D YG hissing sound CaO  H 2 O    Ca(OH )2  Heat, slaked lime When exposed to oxy-hydrogen flame, it starts emitting light called lime light. CaO is used as basic flux, for removing hardness of water, as a drying agent (for NH 3 gas) for preparing mortar (CaO+ sand +water). Mortar : Mortar used in making buildings is a mixture of lime (CaO) and sand in the ratio 1 : 3 with enough water to make a thick paste. When the mortar is placed between bricks, it slowly absorbs CO from the air and the slaked lime revers to CaCO. Ca(OH)2 (s)  CO 2 (g) CaCO 3 (s)  H 2 O(l) Although the sand in the mortar is chemically inert, the grains are bound together by the particles of calcium carbonate and a hard material results. (ii) Calcium chloride (CaCl 2.6 H 2 O) : Fused CaCl 2 is a good dessicant (drying agent). It can't be used to dry alcohol or ammonia as it forms additional products with them. (iii) Calcium carbonate (CaCO ) : 2 ST U 3 3 Ca(OH )2  CO 2 CaCO 3  H 2 O. It is insoluble in water but dissolves in the presence of CO 2 due to the formation of calcium bicarbonate. CaCO 3  H 2 O  CO 2 Ca(HCO 3 )2 It is a constituent of protective shells of marine animals. (iv) Gypsum (CaSO 4. 2 H 2 O) : On partially dehydrates to produce plaster of paris. The above compounds are provided by the two raw materials, namely lime stone (which provides CaO ) and clay which provides SiO 2 , Al 2 O 3 and Fe 2 O 3. In cement, almost entire amount of lime is present in the combined state as calcium silicates (2CaO.SiO2 and 3CaO.SiO 2 ) and calcium aluminates (3CaO. Al 2 O 3 and 4 CaO. Al 2 O 3 ). (b) Cement containing excess amount of lime cracks during setting; while cement containing less amount of lime is weak in strength. (c) Cement with excess of silica is slow-setting and that having an excess of alumina is quick-setting. (d) Cement containing no iron oxide is white but hard to burn. Cement is manufactured by two processes, viz, wet and dry. A small amount (2–3%) of gypsum is added to slow down the setting of the cement so that it gets sufficiently hardened. Setting of cement is an exothermic process and involves hydration of calcium aluminates and calcium silicates. Boron Family Group 13 of long form of periodic table (previously reported as group III A according to Mendeleefs periodic table) includes boron ( B) ; aluminium (Al) , gallium (Ga), indium (In) and thallium (Tl) Boron is the first member of group 13 of the periodic table and is the only non-metal of this group. The all other members are metals. The non-metallic nature of boron is due its small size and high ionisation energy. The members of this family are collectively known as boron family and sometimes as aluminium family. (1) Electronic configuration more effectively pulled the nucleus. This results in less availability of ns electrons pair for bonding or ns electron pair becomes inert. The inert pair effect begins after n  4 and increases with increasing value of n. (iv) The tendency to form M ion increases down the gp. Ga < Tl (10) Hydrated ions : All metal ions exist in hydrated state. (11) Ionisation energy (i) Inspite of the more charge in nucleus and small size, the first ionisation energies of this group elements are lesser than the corresponding elements of s block. This is due to the fact that removal of electron from a p-orbitals (being far away from nucleus and thus less effectively held than sorbitals) is relatively easier than s-orbitals. (ii) The ionisation energy of this group element decrease down the group due to increases in size like other group elements. (iii) However, ionisation energy of Ga are higher than that of Al because of smaller atomic size of Ga due to less effective shielding of 3d electrons in Ga. Thus valence shell exert more effective nuclear charge in Ga to show higher ionisation energies. (12) Electropositive character (i) Electropositive character increases from B to Tl. (ii) Boron is semi metal, more closer to non-metallic nature whereas rest all members are pure metals. (iii) Furthermore, these elements are less electropositive than s-block elements because of smaller size and higher ionisation energies. (13) Oxidation potential (i) The standard oxidation potentials of these element are quite high and are given below, 2 2 Electronic configuration ( ns 2 np 1 ) Element [He] 2 s 2 p Al [ Ne ] 3 s 2 3 p 1 31 Ga 49 1 [ Ar] 3d 10 4 s 2 4 p 1 In [Kr] 4 d 10 5 s 2 5 p 1 81 Tl [ Xe] 4 f 14 5 d 10 6 s 2 6 p 1 (2) Occurrence : The important of this group elements are given below, Boron : Borax (Tincal) (NaBO.10HO), Colemanite (CaBO5HO) 2 4 7 2 2 6 1 2 Boracite (2MgBO.MgCl), Boronatro calcite (CaBO.NaBO.8H O), 3 8 51 2 4 Kernite (Na B O 4H O), Boric acid 2 4 7. 7 2 2 (H BO ) 2 3 3 Aluminium : Corundum (Al O ), Diaspore (Al O.H O), Bauxite (Al O 2 3 2 3 2 2 3. 2H O), and Cryolite (Na AlF ). 2 3 6 ID Physical properties (1) A regular increasing trend in density down the group is due to increase in size. (3) Boron has very high melting point because it exist as giant covalent polymer in both solid and liquid state. (4) Low melting point of Ga (29.8 C) is due to the fact that consists of only Ga molecule; it exist as liquid upto 2000 C and hence used in high temperature thermometry. D YG 0 0 2 E0op for M M3++ 3e E0op for M M+ + e – U (2) Melting points do not vary regularly and decrease from B to Ga and then increase. (5) Boiling point of these elements however show a regular decrease down the group. (6) The abrupt increase in the atomic radius of Al is due to greater screening effect in Al (it has 8 electrons in its penultimate shell) than in B (it has 2 electrons in its penultimate shell) U (7) The atomic radii of group 13 elements are smaller than the corresponding s-block elements. This is due to the fact that when we move along the period, the new incoming electron occupy the same shell whereas the nuclear charge increases regularly showing more effective pull of nucleus towards shell electrons. This ultimately reduces the atomic size. ST (8) The atomic radius of Ga is slightly lesser than of Al because in going from Al to Ga, the electrons have already occupied 3d sub shell in Ga. The screening effect of these intervening electrons being poor and has less influence to decrease the effective nuclear charge, therefore the electrons in Ga experience more forces of attractions towards nucleus to result in lower size of Ga than Al (9) Oxidation state +1 B Al – +1.66 +0.56 +0.34 – +0.18 +0.34 +0.55 Ga In Tl +1.26 (ii) However Boron does not form positive ions in aqueous solution and has very low oxidation potential. (iii) The higher values of standard oxidation potentials are due to higher heats of hydration on account of smaller size of trivalent cations. (iv) Aluminium is a strong reducing agent and can reduce oxides which are not reduced even by carbon. This is due to lower ionisation energy of aluminium than carbon. The reducing character of these elements is Al > Ga > In > Tl. (14) Complex formation : On account of their smaller size and more effective nuclear charge as well as vacant orbitals to accept elements, these elements have more tendency to form complexes than-s block elements. Chemical properties (1) Hydrides (i) Elements of group 13 do not react directly with hydrogen but a number of polymeric hydrides are known to exist. (ii) Boron forms a large no. of volatile covalent hydrides, known as boranes e.g. B H ,B H ,B H ,B H Two series of borones with general formula B H and B H are more important. (iii) Boranes are electron deficient compounds. It is important to note that although BX are well known, BH is not known. This is due of the fact that hydrogen atoms in BH have no free electrons to form p–p back bonding and thus boron has incomplete octet and hence BH molecules dimerise to form B H having covalent and three centre bonds. (iv) Al forms only one polymeric hydride (AlH ) commonly known as alane It contains Al…..H……Al bridges. (v) Al and Ga forms anionic hydrides e.g. LiAlH and LiGa H , 2 n n+4 n 6 4 10 5 11 6 10 n+6 3 3 3 3 (i) All exhibit +3 oxidation state and thus complete their octet either by covalent or ionic union. (ii) Boron being smaller in size cannot lose its valence electrons to form B ion and it usually show +3 covalence. The tendency to show +3 covalence however decreases down the group even Al shows +3 covalence in most of its compounds. (iii) Lower elements also show +1 ionic state e.g Tl , Ga. This is due to inert pair effect. The phenomenon in which outer shell ‘s’ electrons (ns ) penetrate to (n-1) d-electrons and thus become closer to nucleus and are 3+ + + 2 +1 E3 13 2 60 + 5B 6 6 3 n 4 ether 4 LiH  AlCl3   Li[ AlH4 ]  3 LiCl (2) Reactivity towards air 4 3 2 3 2 2 2 3 3 3 3 2 2 3 3 3 3 3 3 2 3 3 2 3 3 3 3 B(OH) + H O B(OH) +H (iv) Al O being amphoteric dissolves in acid and alkalies both. + 1– 2 2 4 3 Al O + 3H SO Al (SO ) + 3H O 2 3 2 4 2 4 3 2 Al2 O 3  2 NaOH  fuse 2 NaAlO3  Sodium meta aluminate H 2O 2 D YG 3 3 (5) Action of Alkalies (i) Boron dissolves only in fused alkalies, 2B + 6NaOH (fused) 2Na BO + 3H (ii) Al and Ga dissolves in fused as well as in aqueous alkalies, 2Al + 2 NaOH + 2H O 2NaAl O + 3H (iii) Indium remains unaffected in alkalies even on heating. 3 2 U 3 2 2 2 ST (6) Halides (i) All the group 13 elements from the trihalides, MX on directly combining with halogens. 3 M + X MX (ii) All the trihalides of group 13 elements are known except Tl (III) 2 3 iodide. 3 BX + 3H O B(OH) + 3HX; [X=Cl, Br, I] However, BF forms as addition product with water, HO BF + H OH [BF OH] H O [BF OH]. 2 3 3 + 2 3 3 3 3 3 - 3 3 3 3 3 3 3 3. – 4 3– 3– 6 3– 6 6 6 0 2 3 3 2 (iii) Due to small size and high electronegativity of boron, all boron halides are covalent and Lewis acids. These exist as monomeric molecules having plane triangular geometry (sp hybridization). (iv) All Boron trihalides except BF are hydrolysed to boric acid. 3 3 U (v) One of the crystalline form of alumina (Al O ) is called corrundum. It is very hard and used as abrasive. It is prepared by heating amorphous form of Al O to 2000 K. (4) Action of Acids (i) Boron does not react with non oxidizing acids, however, it dissolves in nitric acid to form boric acids. (ii) Al, Ga and In dissolve in acids forming their trivalent cations; however, Al and Ga become passive due to the formation of protective film of oxides. (iii) Thallium dissolves in acids forming univalent cation and becomes passive in HCl due to the formation of water insoluble TICl. 2 3 ID 3 BF having less tendency for hydrolysis as well as Lewis acid nature, is extensively used as a catalyst in organic reactions e.g. Friedel- Crafts reaction. (v) Boron atom, in BX , has six electrons in the outermost orbit and thus it can accept a pair of electrons form a donor molecule like NH to complete its octet. Hence boron halides act as very efficient Lewis acids. The relative Lewis acid character of boron trihalides is found to obey the order ; BI >BBr >BCl >BF. However, the above order is just the reverse of normally expected order on the basis relative electronegativities of the halogens. Fluorine, being the most electronegative, should create the greatest electron deficiency on boron and thus B in BF should accept electron pair from a donor very rapidly than in other boron trihalides. But this is not true. This anomalous behaviour has been explained on the basis of the relative tendency of the halogen atom to back-donate its unutilised electrons to the vacant p orbitals of boron atom. In boron trifluoride, each fluorine has completely filled unutilised 2p orbitals while boron has a vacant 2p orbital. Now since both of these orbitals belong to same energy level (2p) they can overlap effectively as a result of which fluorine electrons are transferred into the vacant 2p orbital of boron resulting in the formation of an additional p–p bond. This type of bond formation is known as back bonding or back donation. Thus the B- F bond has some double bond character. Back bonding may take place between boron and of the three fluorine atoms and thus boron trifluoride is regarded as a resonance hybrid of some structures. Resonance in boron trifluoride is also evidenced by the fact that the three boron-fluorine bonds are indentical and are shorter than the usual single boron-fluorine bond As a result of back bonding, the electron deficiency of boron is reduced and hence Lewis acid nature is decreased. The tendency for the formation of back bonding (p- p bond) is maximum in BF and decreases very rapidly from BF to BI This is probably due to the fact that overlapping of the vacant 2p orbitals of boron cannot take place easily with the p-orbitals of high energy levels (3p in Cl, 4p in Br and 5p in iodine). Thus BI Br and BCl are stronger Lewis acids than the BF (vi) Lewis acid character of halides of the group 13 elements decreases in the order, B > Al > Ga > In. (vii) Boron halides form complex halides of the type, [BF ], in which boron atom extends its coordination number to four by utilising empty porbital. It cannot extend its coordination number beyond four due to non availability of d-orbitals. However, the other trihalides of this group form complex halides of the type (AlF ) , (GaCl ) and (InCl ) , etc where the central atom extends its coordination number to 6 by the use of d-orbitals. (viii) The fluorides of Al, Ga In and Tl are ionic and have high melting points. The high melting points of metal fluorides can be explained on the basis that their cations are sufficiently large and have vacant dorbitals for attaining a coordination number of six towards the relatively small fluorine atom. (ix) Other halides of Al, Ga, In and Tl are largely covalent in anhydrous state and possess low melting point. These halides do not show backbonding because of increases in the size of the element. However, the make use of vacant p-orbitals by co-ordinate bond i.e. metal atoms complete their octet by forming dimers. Thus aluminium chloride, aluminium bromide and indium iodide exist as dimers, both in the vapour state and in non-polar solvents. The dimer structure for Al Cl is evidenced by the following facts, (a) Vapour density of aluminium chloride measured at 400 C corresponds to the formula Al Cl. (b) Bond distance between aluminium chlorine bond forming bridge is greater (2.21Å) than the distance between aluminum-chlorine bond present in the end (2.06 Å). The dimeric structure disappears when the halides are dissolved in water This is due to high heat of hydration which split the dimeric structure into [ M(H O) ] and 3X ions and the solution becomes good conductor of electricity. 60 2 E3 (i) Pure boron is almost unreactive at ordinary temperature. It reacts with air to form B O when heated It does react with water. Al burns in air with evolution of heat give Al O. (ii) Ga and In are not effected by air even when heated whereas Tl is little more reactive and also form an oxide film at surface. In moist air, a layer of Tl (OH) is formed. (iii) Al decomposes H O and reacts readily in air at ordinary temperature to form a protective film of its oxides which protects it from further action. (3) Oxides and hydroxides (i) The members of boron family form oxide and hydroxides of the general formula M O and M (OH) respectively. (ii) The acidic nature of oxides and hydroxides changes from acidic to basic through amphoteric from B to Tl. B O and B(OH) > Al O and Al(OH) > (acidic) (amphoteric) Ga O and Ga(OH) > In O In (OH) > Tl O Tl(OH) (amphoteric) (basic) (strong basic) B(OH) or H BO is weak monobasic Lewis acid. (iii) Boric acid, B(OH) is soluble in water as it accepts lone pair of electron to act as Lewis acid. Rest all hydroxides of group 13 are insoluble in water and form a gelatinous precipitate. + 2 3 - 3 2 6 3+ 2 – 6 Al Cl + 2H O 2[Al(H O) ] +6Cl ; Therefore Al Cl is ionic in water. 3+ 2 6 2 2 6 – 2 6 3 3 2 6 3 3 3 6 D YG 2 – 4 3- 6 3 2 B + 3HNO H BO + 3NO 3 3 3 2 ST U Diagonal relationship between Boron and Silicon Due to its small size and similar charge/mass ratio, boron differs from other group 13 members, but it resembles closely with silicon, the second element of group 14 to exhibit diagonal relationship. Some important similarities between boron and silicon are given below, (1) Both boron and silicon are typical non-metals, having high m.pt. b.pt nearly same densities (B=2.35gml S=2.34 g//ml). low atomic volumes and bad conductor of current. However both are used as semiconductors. (2) Both of them do not form cation and form only covalent compounds. (3) Both exists in amorphous and crystalline state and exhibit allotropy. (4) Both possess closer electronegativity values (B=2.0; Si=1.8). (5) Both form numerous volatile hydrides which spontaneously catch fire on exposure to air and are easily hydrolysed. (6) The chlorides of both are liquid, fume in most air and readily hydrolysed by water. BCl + 3H O B(OH) + 3HCl SiCl + H O Si(OH) + 4HCl 3 2 4 2 3 4 –1 3 2 3 3 4 3Mg+2BMg B ; Mg B +H PO Mixture of boranes 3 2 3 2 3 4 (Magnesium boride) 2Mg + Si Mg Si ; Mg Si + H PO Mixture of silanes 2 2 3 4 (magnesium silicide) (9) The carbides of both Boron and silicon (B C and SiC) are very hard and used as abrasive. (10) Oxides of both are acidic and can be reduced by limited amount of Mg In excess of Mg boride and silicide are formed. 4 B O +3Mg 3MgO+2B ; SiO +2Mg 2MgO+Si (11) Both the metals and their oxides are readily soluble in alkalies. 2B + 6NaOH 2Na BO + 3H  2 3 2 3 3 2 (borate) Si + 2NaOH + H O Na SiO + 2H  2 2 3 2 (silicate) B O + 6NaOH 2Na BO + 3H O SiO + 2NaOH Na SiO + H O 2 3 3 2 3 2 2 3 2 Both borates and silicates have tetrahedral structural units SiO4n  respectively. Boro silicates are known in which boron replaces silicon in the three dimensional lattice. Boron can however form planar BO units. (12) Acids of both these elements form volatile esters on heating with alcohol in presence of conc. H SO. B(OH) + 3ROHB(OR) + 3H O Si(OH) + 4ROH Si(OR) + 4H O BO4n  and 3 2 3 U 2 3 ID 3 (7) Both form weak acids like H BO and H SiO. (8) Both form binary compounds with several metals to give borides and silicide. These borides and silicide react with H PO to give mixture of boranes and silanes. 60 3 E3 The dimeric structure may also split by reaction with donor molecules e.g. R N. This is due to the formation of complexes of the type R NAlCl The dimeric structure of Al Cl exist in vapour state below 473K and at higher temperature it dissociates to trigonal planar AlCl molecule. Boron halides do not exist as dimer due to small size of boron atom which makes it unable to co-ordinate four large-sized halide ions. (x) BF and AlCl acts as catalyst and Lewis acid in many of the industrial process. Anomalous Behaviour of Boron Like Li and Be, Boron – the first member of group 13 also shows anomalous behaviour due to extremely low size and high nuclear charge/size ratio, high electronegativity and non-availability of d electrons. The main point of differences are, (1) Boron is a typical non- metal whereas other members are metals. (2) Boron is a bad conductor of electricity whereas other metals are good conductors. (3) Boron shows allotropy and exists in two forms – crystalline and amorphous. Aluminium is a soft metal and does not exist in different forms. (4) Like other non-metals, the melting point and boiling point of boron are much higher than those of other elements of group 13. (5) Boron forms only covalent compounds whereas aluminium and other elements of group 13 form even some ionic compounds. (6) The hydroxides and oxides of boron are acidic in nature whereas those of others are amphoteric and basic. (7) The trihalides of boron (BX ) exist as monomers On the other hand, aluminium halides exist as dimers (Al X ). (8) The hydrides of boron i.e. boranes are quite stable while those of aluminium are unstable. (9) Dilute acids have no action on boron Others liberate H from them. (10) Borates are more stable than aluminates. (11) Boron exhibit maximum covalency of four e.g., BH ion while other members exhibit a maximum covalency of six e.g., [Al(OH) ]. (12) Boron does not decompose steam while other members do so. (13) Boron combines with metals to give borides e.g. Mg B. Other members form simply alloys. (14) Concentrated nitric acid oxidises boron to boric acid but no such action is noticed other group members. 4 3 2 4 4 2 Boron and its compounds Boron is the first member of group –13 (IIIA) of the periodic table. Boron is a non- metal. It has a small size and high ionization energy due to which it can not lose its valence electrons to form B 3 ion. Its compounds especially the hydrides and halides are electron deficient and behave as Lewis acid. (1) Ores of boron (i) Borax or tincal : Na B O. 10H O (ii) Kernite or Rasorite : Na B O. 4H O (iii) Colemanite : Ca B O. 5H O (iv) Orthoboric acid : H BO (It occurs in the jets of steam called soffioni escaping from ground in the volcanic region of the Tuscany). Boron is present to a very small extent (0.001%) in earth’s crust. (2) Isolation : Elemental boron in the form of dark brown powder is obtained either by reduction of boric oxide with highly electropositive metals like K, Mg, Al, Na, etc. in the absence of air and boron halides with hydrogen at high temperature eg. 2 4 7 2 2 2 6 4 11 7 2 2 3 3 Heat  2B + 3K O B O + 6K  2 3 2 2BCl + 3H   2B + 6HCl. By thermal decomposition of boron triiodide over red hot tungsten filament and boron hydrides for example, 1270 K 3 2 W, heat Heat 2BI    2B + 3I ; B H   2B + 3H (3) Properties : It exists in mainly two allotropic forms i.e. amorphous dark brown powder and crystalline black very hard solid. It occurs in two isotopic forms, i.e., 5 B10 (20% abundance) and 5 B11 (80% 3 2 2 6 2 abundance). With air, boron forms B2 O3 and BN at 973K, with halogens, trihalides (BX 3 ) are formed, with metals borides are formed. eg. Heat 4B+ 3O 2   Heat  2B + N 2  2 B2 O3 Boron trioxide 2BN Boron nitride 2B + 3 X 2  2BX 3 Boron trihalide Heat 3Mg + 2B   H Mg3 B2 Magnesium boride H.... t or B Water, steam and HCl have no action on B. oxidising acids (HNO 3 , H 2 SO 4 ) convert boron to H 3 BO3. H B + 3 HNO 3  H 3 BO3 + 3 NO 2 H t H b t B H b t Structure of diborane (B H ) 2 2B + 3 H 2 SO 4  2 H 3 BO3 + 3 SO 2 60 The other boron hydrides are B5 H 9 , B4 H10 , B5 H11 etc. (ii) Boron Halides Boron reacts with halogens on strong heating to form boron halides Fused alkalies (NaOH, KOH) dissolve boron forming borates, liberating hydrogen. Fused 2B + 6KOH    2 K3 BO3 + 3H 2. boranes. Diborane is the simplest boron hydride which is a dimer of BH 3. 450 K (a) 8 BF3  6 LiH   B2 H 6  6 LiBF4 (b) 4 BCl 3  LiAlH4  2 B2 H 6  3 AlCl3  3 LiCl In these halides, the central boron atom has three shared pairs of electrons with the halogen atoms. Therefore, these have two electrons less than the octet and are electron deficient compounds. They acts as Lewis acids. H F H F | | | | F  B  : N  H  F  B  N  H | | | | F H F H Lewis acid Lewis base The relative acidic strength of boron trihalides decreases as : BI3  BBr3  BCl 3  BF3. U (c) In the laboratory, it is prepared by the oxidation of sod. Borohydride with I2. Polyether 2 NaBH 4  I2   B2 H 6  2 NaI  H 2 BF3 and BCl 3 are gases, BBr3 is a volatile liquid while BI3 is a solid. ID Preparation Heat 2B + 3 X 2   2B X 3 (X  F, Cl, Br, I) E3 (4) Uses of Boron : Boron is used in atomic reactors as protective shields and control rods, as a semiconductors for making electronic devices in steel industry for increasing the hardness of steel and in making light composite materials for air crafts. (5) Compounds of Boron (i) Boron Hydrides Boron forms hydrides of the types Bn H n  4 and Bn H n  6 called 6 D YG Properties : (a) Since Boron in boranes never complete its octet of electrons hence all boranes are called as electron-deficient compounds or Lewis acids. (b) All boranes catch fire in the presence of oxygen to liberated a lot of heat energy. Thus, they can also be used as high energy fuels. (iii) Borax ( Na 2 B 4 O7. 10 H 2 O ) It occurs naturally as tincal (Suhaga) which contains about 50% borax in certain land, lakes. It is also obtained from the mineral colemanite by boiling it with a solution of Na 2 CO 3. Ca 2 B6 O11  2 Na 2 CO 3  Na 2 B4 O7  2CaCO 3  2 NaBO2 Colemanite Borax B2 H 6  3O2  2 B2 O3  3 H 2 O; H  1976 KJ / mole Properties : (a) Its aqueous solution is alkaline due to hydrolysis, (c) Boranes are readily hydrolysed by water. Na 2 B4 O7 + 7 H 2 O B2 H 6  6 H 2 O  2 H 3 BO3  6 H 2 (d) With carbon monoxide B2 H 6  2CO (BH 3  CO )2 U (e) Boranes are used for formation of hydroborates or borohydrides such as LiBH 4 or NaBH 4 , which are extensively used as reducing agents in organic synthesis. Diethyl ether 2 LiH  B2 H 6    2 Li  [BH 4 ] ST Structure of diborane : B2 H 6 has a three centre electon pair bond also called a banana shape bond. (a) B  H t : It is a normal covalent bond (two centre electron pair bond i.e., 2c  2e ). (b) B  H b : This is a bond between three atoms, B  H b  B, (three centre electron pair bond i.e., 3c  2e ). H t B H t 97 H b B o H b   Na 2 B 4 O7.10 H 2 O  Na 2 B 4 O7  2 NaBO2  B 2 O3 10 H 2O Borax bead Borax bead is used for the detection of coloured basic radicals under B 2 O 3 on heating combines readily with a number of coloured transition metal oxides such as Co, Ni, Cr, Cu, Mn, etc. to form the corresponding metaborates which possess characteristic colours,  CoSO 4  CoO  SO 3 ; CoO  B 2 O 3  Co (BO 2 )2 Cobalt meta borate (Blue) Colours of some important metaborates are : Cupric metaborate, Cu (BO 2 )2 is dark blue, chromium metaborate, Cr (BO2 )2 is green, nickel metaborate, Ni(BO2 )2 is brown and manganese metaborate, Mn(BO2 )2 is pink violet. (c) When heated with C 2 H 5 OH and conc. H 2 SO 4 it gives volatile vapours of triethyl borate which burns with a green edged flame. Bridged Hydrogen H (b) On heating borax loses its water of crystallization and swells up to form a fluffy mass. On further heating, it melts to give a clear liquid which solidifies to a transparent glassy bead consisting of sodium metaborate ( NaBO2 ) and boric anhydride ( B2 O3 ), the name borax bead test. Borax bead test : Borax bead is a mixture of NaBO2 and B 2 O 3. Diethyl ether 2 NaH  B2 H 6   2 Na  [BH 4 ] Terminal Hydrogen  2NaOH+4 H 3 BO3. t 120 Na 2 B4 O7  H 2 SO 4  5 H 2 O  Na 2 SO 4  4 H 3 BO3 o H t It is obtained from borax by treating with dil. HCl or dil. H 2 SO 4 , Na 2 B 4 O7  2 HCl  5 H 2 O  2 NaCl  4 H 3 BO3 It can also be obtained from the mineral colemanite by passing through a mixture of powdered mineral in boiling water, Ca 2 B6 O11 + 4 SO 2  11H 2 O  2Ca(HSO 3 )2  6 H 3 BO3 Properties : (a) It is a very weak monobasic acid, does not act as a proton doner but behaves as a Lewis acid i.e. it accepts a pair of electrons from OH  ion of H 2 O , H 3 BO3  H 2 O [B(OH )4 ]  H  It acts as a strong acid in presence of polyhydroxy compounds such as glycerol, mannitol etc. and can be titrated against strong alkali. (b) With NaOH it forms, sodium metaborate, Boron oxide Tetra boric acid BO 33 D YG Structure : In boric acid, planar units are joined by hydrogen bonds to give a layer structure. Uses : (a) As a food preservative. (b) As a mild antiseptic for eye wash under the name boric lotion. (c) For the preparation of glazes and enamels in pottery. (v) Borazine or Borasole or Triborine triamine ( B3 N 3 H 6 ) It is a compound of B, N and H. It is a colourless liquid and is also called inorganic benzene. o 180 C 2 B2 H 6  6 NH 3    2 B3 N 3 H 6  12 H 2. It has a six membered ring of alternating B and N atoms, each is further linked to a H- atom. FeO Fe2 O3 High pressure (150 o C, 80 atm) filtered,Fe 2 O3 as residue Filtered Heat Filtrate   Pure Al2 O3   Al(OH )3 (Sod. Aluminate) CO 2 (ii) Hall's process Bauxite CO 3  Na 2   Solution 2 CO o (Finely powdered) Fused, extracted with water. Residue Fe 2 O3 (red) 50 - 60 C and filtered. Filtrate(Na 2 CO 3 ) Heat Precipitate Al(OH)3   Pure Al2 O3 (iii) Serpek's process  Coke  N Bauxite (Finely powdered) (white) 2   Silica reduced to  Alumina form AIN Heated to Si which volatalises aluminium nitride o 1800 C (iv) Hall and Heroult process : It is used for extraction of aluminium. In this process a fused mixture of alumina (20%), cryolite (60%) and fluorspar (20%) is electrolysed using carbon electrodes. Whereas cryolite makes Al2 O3 conducting fluorspar decreases the melting point of alumina. Aluminium is refined by Hoope's electrolytic process. (3) Compounds of Aluminium (i) Aluminium oxide or Alumina ( Al2 O3 ) : It occurs in nature as colourless corundum and several coloured minerals like ruby (red), topaz (yellow), Sapphire (blue), amethyst (violet) and emerald (green). These minerals are used as precious stones (gems). (ii) Aluminium chloride ( Al2 Cl 6 ) : It is prepared by passing dry chlorine over aluminium powder. Al2 O3  3 C  3 Cl 2 2 AlCl3  3 CO (g) (anhydrous) It exists as dimer Al2 Cl 6 , in inert organic solvents and in vapour B when AlCl3 is dissolved in water. It is hygroscopic in nature and absorbs moisture when exposed to air. (iii) Thermite : A mixture of aluminium powder and Fe2 O3 in the N—H U ST Roasted  Caustic soda solution   Roasted ore   state. It sublimes at 100 o C under vacuum. Dimeric structure disappears B—H ratio 1:3. It is used for welding of iron. The reaction between Al and Fe2 O3 is highly exothermic, N H Al  Fe2 O3 Al2 O3  Fe  Heat Borazine It is prepared by treating BCl 3 with an excess of NH and pyrolysing the resulting mixture in an atmosphere of NH at 750°C, 3 3 750 o C  BN + 3HCl. BCl 3 + NH 3  [ H 3 N  BCl 3 ]  Excess NH 3 It is a colourless, good insulator, diamagnetic and almost unreactive solid Aluminium and its compounds (1) Ores of Aluminium : Bauxite ( Al2 O 3.2 H 2 O), Cryolite (Na 3 AlF6 , Felspar (KAlSi3 O8 ), Kaolinite ( Al2 O3. 2SiO2.2 H 2 O) , Mica (K 2 O.3 Al2 O3. 6 SiO2.2 H 2 O), Bauxite Finely powdered (red) H H—N (vi) Boron nitride (BN) (2) Extraction : Aluminium is obtained by the electrolysis of the oxide (alumina) dissolved in fused cryolite. This involves following steps, Purification of ore (i) Baeyer's process U 433 K Red hot    H 2 B4 O7   B2 O3 H—B [K2 SO 4. Al2 (SO 4 )3. ID  B(O C 2 H 5 )3 + 3 H 2 O H 3 BO3 +3 C2 H 5 OH  (d) Action of heat : The complete action of heat on boric acid may be written as, Conc. H 2 SO 4 Metaboric acid stone Heated   Pure Al2O3    Al(OH )3 (c) With C 2 H 5 OH and conc. H 2 SO 4 , it gives triethyl borate Boric acid alum Hydrolysis H 3 BO3  NaOH  NaBO2  2 H 2 O 373 K H 3 BO3   HBO2 or 4 Al(OH )3 ]. E3 This reaction is used as a test for borate radical in qualitative analysis. Uses : (a) In making optical and hard glasses. (b) In the laboratory for borax bead test. (c) In softening of water. (d) In the preparation of medicinal soaps due to its antiseptic character. (iv) Boric acid or orthoboric acid ( H 3 BO 3 ) SO 2 Alunite ( Al2 O3.H 2 O), Triethyl borate 60 H 3 BO3  3 C 2 H 5 OH  B(OC2 H 5 )3  3 H 2 O Corundum ( Al2 O3 ) , Diaspore (iv) Aluminium sulphate [Al (SO ) ] : It is used for the preparation of alums e.g., potash alum Al2 (SO 4 )3. K 2 SO 4. 24 H 2 O. It is also used for making fire proof clothes. (iv) Alums : In general, the term alum is given to double sulphates of the type M 2 SO 4.M 2 (SO 4 )3.24 H 2 O where M is a univalent 2 4 3 cation like Na  , K  and NH 4 , M  is a trivalent cation like Al 3  , Fe 3  and Cr 3 . Some important points to be noted about the alums are (a) General formula is M 2 SO 4.M 2 (SO 4 )3.24 H 2 O M  Monovalent metal; M   Trivalent metal In alum crystals, 6 water molecules are held by monovalent ion, 6 water molecules are held by trivalent ion, 12 water molecules are held in the crystal structure. (b) All alums are isomorphous. Aqueous solutions of alums are acidic due to cationic hydrolysis of trivalent cation. (c) Double sulphates of divalent ions and trivalent ions with 24 water molecules in their crystals are known as Pseudoalums. General formula is MSO 4.M 2 (SO 4 )3.24 H 2 O M  Bivalent metal; M   Trivalent metal (d) Pseudoalums are not isomorphous with alums. E3 Potash alum K 2 SO 4. Al 2 (SO 4 )3.24 H 2 O Sodium alum Na 2 SO 4. Al 2 (SO 4 )3.24 H 2 O Ammonium alum (NH 4 ) 2 SO 4. Al 2 (SO 4 )3.24 H 2 O Chrome alum K 2 SO 4.Cr2 (SO 4 )3.24 H 2 O ID Carbon Family 14 Si 32 Ge Sn 82 Pb [He] 2 s 2 2 p 2 [ Ne ] 3 s 2 3 p 2 [ Ar] 3d 10 4 s 2 4 p 2 [Kr] 4 d 10 5 s 2 5 p 2 [ Xe] 4 f 14 5 d 10 6 s 2 6 p 2 ST U 50 Electronic configuration ( ns 2 np 2 ) D YG 6C U Carbon is the first member of group 14 or IVA of the periodic table. It consists of five elements carbon (C), silicon (Si), germanium (Ge), tin (Sn) and lead (Pb). Carbon and silicon are nonmetals, germanium is metalloid and tin and lead are metals. (1) Electronic configuration Elements 60 (e) Feather alum or ‘Hair-salt’ Al 2 SO 4.18 H 2 O is a native form of aluminium sulphate. (f) Potash alum is used for tanning of leather, as mordant in dyeing and calico printing, for sizing paper, as a syptic to stop bleeding and purification of water. Some important alums are Physical properties (1) Non-metallic nature : The non-metallic nature decreases along the group. C Si Ge Non-metals metalloid The electronegativity from silicon onwards is almost is almost constant or shows a comparatively smaller decreases due to screening effects of d electrons in elements from Ge onwards. 10 (7) Ionisation energy Sn Pb metal metal (2) Abundance : Carbon and silicon are most abundant elements in earth’s crust whereas germanium occurs only as traces. Tin and lead also occur in small amounts. Only carbon occurs in free state as coal, diamond and graphite and in combined state as carbonates, CO petroleum and natural gas Silicon is the second most abundant element after oxygen in earth’s crust in form of silicates and silica. Germanium found in traces in coal and in certain deposits. It important constituent for making conductors and transistors The important ore of tin is tin stone (SnO ) or cassiterite. Lead is found is form of galena (PbS) anglesite (PbSO ) and cerussite (PbCO ) The abundance ratio in earth’s crust is given below, (i) The ionisation energy decreases regularly down the group; Pb however shows a higher value than Sn due to poor shielding of inner forbitals as a result of which effective nuclear charge experienced by outer shell electrons becomes more in Pb. Ionisation energy (kJ mol ) -1 IE 4 3 IE C Si Density (g/ml) Ge 3.51 (for diamond) 2.34 Sn Pb 5.32 7.26 11.34 (i) The melting point and boiling point of this group members decrease down the group. C Si Ge Sn m.pt(K) 4373 1693 1218 505 b.pt.(K) – 3550 3123 2896 Pb D YG U Ge Sn Pb 0.77 111 122 141 144 Atomic volume (ml) 3.4 11.4 13.6 16.3 18.27 ST 1411 1450 4- 4+ (ii) The formation of M or M ions require huge amount of energy which is normally not available during normal course of reactions, therefore, these elements usually do not form M or M ions, but they usually form compounds with covalence of four. 4+ 4- 4+ 4- (iii) Some of the ionic radii involving six co-ordination of these group elements are given below, Sn C Si Ge Ionic radius (M ) in pm – – 73 118 119 Ionic radius (M ) in pm – 40 53 69 78 Pb (6) Electronegativity : The electronegativity decreases from C to Si and then becomes constant. C Si Ge Sn Pb 2.5 1.8 1.8 1.7 1.6 2+ 2+ (v) The tendency to form +2 ionic state increases on moving down the group due to inert pair effect. (9) Catenation (i) The tendency of formation of long open or closed atom chains by the combination of same atoms in themselves is known as catenation. (ii) The catenation is maximum in carbon and decreases down the group. (iii) This is due to high bond energy of catenation. (iv) Only carbon atoms also form double or triple bonds involving p-p multiple bond within itself. > C = C< ; – C  C – (ii) The atomic radii of group 14 elements are than their corresponding group 13 elements due to increase in nuclear charge in the same period. Electronegativity on pauling scale 1537 (iv) Sn and Pb show ionic nature. Atomic radius (pm) ++ 1577 (i) Presence of four electrons in outermost shell of these elements reveals that the members of this family can gain four electrons forming M or M ions to show ionic nature or exhibit tetravalent covalent nature by sharing of four electron pairs in order to attain stable configuration. 2024 (i) Both atomic radii and atomic volume increases gradually on moving down the group due to the effect of extra shell being added from member to member. 2+ 2352 2 (iii) Ge, Sn and Pb also exhibit +2 oxidation state due to inert pair (5) Atomic radii and atomic volume Si 715 effect. 600 (ii) The melting point and boiling point of group 14 elements are however, higher than their corresponding group 13 elements. This is due to the formation of four covalent bonds on account of four electrons in their valence shells which results in strong binding forces in between their atoms in solid as well as in liquid state. C Pb 708 (8) Oxidation state U Element Sn 761 (iii) The electropositive character of these elements increases down the group because of decreases in ionisation energy. ID 2.22 (for graphite) (4) Melting point and boiling points Ge (ii) The first ionisation energies of group 14 elements are higher than their corresponding group 13 elements because of smaller size. (3) Density : The density of these elements increases down the group as reported below Element Si 786 E3 2 1 C 1086 60 2 (v) Carbon also possesses the tendency to form closed chain compounds with O, S and N atoms as well as forming p-p multiple bonds with other elements particularly nitrogen and oxygen e.g. C =O; C=N; C  N; C = S are the functional groups present in numerous molecules due to this reason. (vi) Carbon can form chain containing any number of carbon atoms Si and Ge cannot extend the chain beyond 6 atoms, while Sn and Pb do not form chains containing more than one or two atoms. (vii) The reason for greater tendency of carbon for catenation than other elements in the group may further be explained by the fact that the C – C bond energy is approximately of the same magnitude as the energies of the bond between C and other elements. On the other hand, the Si – Si bond is weaker than the bond between silicon and other elements. Bond Bond energy (kJ/mol) C–C 348 Si–Si 180 C–O 315 Si–O 372 C–H 414 Si–H 339 C–Cl 326 Si–Cl 360 C–F 439 Si–F 536 (iii) Tin has three crystalline modifications with the following equilibrium temperature   tin 15.2oC (Grey) (10) Allotropy The phenomenon of existence of a chemical element in two or more forms differing in physical properties but having almost same chemical nature is known as allotropy. If an element or compound exists in two or more forms, it is also known as polymorphism e.g. zinc blende and wurtzite are polymorphs of ZnS. Kinds of allotropy. Allotropy is of three types :   Sn 161oC (White) Liquid tin The conversion of white tin to grey tin is accompanied by an increase in volume and the latter, being very brittle, easily crumbles down to powder. This phenomenon is called tin disease tin pest or tin plague. Chemical properties (1) Hydrides : All the elements of group 14 combine with hydrogen directly or indirectly to form the covalent hydrides, MH 4 (M = C, Si, Ge, Sn or Pb). The number of hydrides and the ease of preparation decrease on going from carbon to lead. The hydrides of silicon are called silanes having the general formula Sin H 2n  2. The hydrides of germanium are called germanes while those of tin are called the stannanes. Only lead forms an unstable hydride of the formula, PbH4 called the plumbane. Three hydrides of germanium, i.e., GeH 4 , Ge 2 H 6 and Ge 3 H 8 and only two hydrides of tin i.e., SnH 4 and Sn 2 H 6 are well known. (2) Oxides : Carbon forms five oxides CO, CO 2 , C3 O2 (carbon suboxide), C5 O 2 and C12 O9 , C3 O2 is the anhydride of malonic acid and CO 2 is the anhydride of H 2 CO 3 (carbonic acid) CO 2 is a non-polar ID (i) Enantiotropy : When two forms of a solid substance exist together in equilibrium with each other at a particular temperature under normal pressure it is called enantiotropy. For example, at normal pressure and temperature between 368.6 K and 285 K , sulphur (solid) exist in two forms (rhombic sulphur), S R and (monoclinic sulphur), S M in equilibrium with each linear molecule due to maximum tendency of C to form p–p multiple bond with oxygen. Si forms SiO2. Pb forms a number of oxides. PbO can be obtained U other. S R ⇌ S M   Sn 232oC (Rhombic) 60 Bond energy (k J/mol) E3 Bond D YG (ii) Monotropy : It is the type of allotropy in which only one allotrope is stable, under normal conditions the other being unstable e.g., diamond and graphite, oxygen and ozone etc. (iii) Dynamictropy : It is the type of allotropy in which there is a true equilibrium between the two allotropes, one changing into the other at exactly the same rate as the reverse occurs. Both allotropes are stable over a wide range of