kech203.pdf

THE s-BLOCK ELEMENTS 291 UNIT 10 THE s -BLOCK ELEMENTS...

THE s-BLOCK ELEMENTS 291 UNIT 10 THE s -BLOCK ELEMENTS d he The first element of alkali and alkaline earth metals differs in many respects from the other members of the group is After studying this unit, you will be able to bl The s-block elements of the Periodic Table are those in describe the general charact- which the last electron enters the outermost s-orbital. As eristics of the alkali metals and pu the s-orbital can accommodate only two electrons, two their compounds; groups (1 & 2) belong to the s-block of the Periodic Table. explain the general characteristics Group 1 of the Periodic Table consists of the elements: of the alkaline earth metals and lithium, sodium, potassium, rubidium, caesium and be T their compounds; francium. They are collectively known as the alkali metals. These are so called because they form hydroxides on re describe the manufacture, o R properties and uses of industrially reaction with water which are strongly alkaline in nature. important sodium and calcium The elements of Group 2 include beryllium, magnesium, compounds including Portland calcium, strontium, barium and radium. These elements tt E cement; with the exception of beryllium are commonly known as appreciate the biological the alkaline earth metals. These are so called because their significance of sodium, oxides and hydroxides are alkaline in nature and these C potassium, magnesium and metal oxides are found in the earth’s crust*. calcium. Among the alkali metals sodium and potassium are abundant and lithium, rubidium and caesium have much no N lower abundances (Table 10.1). Francium is highly 223 radioactive; its longest-lived isotope Fr has a half-life of only 21 minutes. Of the alkaline earth metals calcium and magnesium rank fifth and sixth in abundance respectively © in the earth’s crust. Strontium and barium have much lower abundances. Beryllium is rare and radium is the –10 rarest of all comprising only 10 per cent of igneous rocks† (Table 10.2, page 299). The general electronic configuration of s-block elements 1 2 is [noble gas]ns for alkali metals and [noble gas] ns for alkaline earth metals. † A type of rock formed * The thin, rocky outer layer of the Earth is crust. from magma (molten rock) that has cooled and hardened. 291 C:\ChemistryXI\Unit-10\Unit-10-Lay-3(reprint).pmd Reprint 27.7.6  292 CHEMISTRY Lithium and beryllium, the first elements increase in atomic number, the atom becomes + of Group 1 and Group 2 respectively exhibit larger. The monovalent ions (M ) are smaller some properties which are different from those than the parent atom. The atomic and ionic of the other members of the respective group. radii of alkali metals increase on moving down In these anomalous properties they resemble the group i.e., they increase in size while going the second element of the following group. from Li to Cs. Thus, lithium shows similarities to magnesium 10.1.3 Ionization Enthalpy and beryllium to aluminium in many of their The ionization enthalpies of the alkali metals properties. This type of diagonal similarity is d are considerably low and decrease down the commonly referred to as diagonal relationship group from Li to Cs. This is because the effect in the periodic table. The diagonal relationship of increasing size outweighs the increasing he is due to the similarity in ionic sizes and /or nuclear charge, and the outermost electron is charge/radius ratio of the elements. very well screened from the nuclear charge. Monovalent sodium and potassium ions and divalent magnesium and calcium ions are 10.1.4 Hydration Enthalpy found in large proportions in biological fluids. is The hydration enthalpies of alkali metal ions These ions perform important biological decrease with increase in ionic sizes. functions such as maintenance of ion balance + + + + + Li > Na > K > Rb > Cs and nerve impulse conduction. bl + Li has maximum degree of hydration and 10.1 GROUP 1 ELEMENTS: ALKALI for this reason lithium salts are mostly METALS pu hydrated, e.g., LiCl· 2H2O The alkali metals show regular trends in their 10.1.5 Physical Properties physical and chemical properties with the All the alkali metals are silvery white, soft and increasing atomic number. The atomic, be T light metals. Because of the large size, these physical and chemical properties of alkali elements have low density which increases re metals are discussed below. down the group from Li to Cs. However, o R 10.1.1 Electronic Configuration potassium is lighter than sodium. The melting All the alkali metals have one valence electron, and boiling points of the alkali metals are low 1 indicating weak metallic bonding due to the tt E ns (Table 10.1) outside the noble gas core. The loosely held s-electron in the outermost presence of only a single valence electron in valence shell of these elements makes them the them. The alkali metals and their salts impart C most electropositive metals. They readily lose characteristic colour to an oxidizing flame. This + electron to give monovalent M ions. Hence they is because the heat from the flame excites the are never found in free state in nature. outermost orbital electron to a higher energy no N level. When the excited electron comes back to Element Symbol Electronic configuration the ground state, there is emission of radiation in the visible region as given below: Lithium Li 1s22s1 Metal Li Na K Rb Cs Sodium Na 1s22s22p63s1 © Potassium K 1s22s22p63s23p64s1 Colour Crimson Yellow Violet Red Blue Rubidium Rb 1s 2s 2p 3s 3p 3d 4s 4p 5s 2 2 6 2 6 10 2 6 1 red violet Caesium Cs 1s22s22p63s23p63d104s2 λ/nm 670.8 589.2 766.5 780.0 455.5 4p64d105s25p66s1 or [Xe] 6s1 Alkali metals can therefore, be detected by Francium Fr [Rn]7s1 the respective flame tests and can be determined by flame photometry or atomic 10.1.2 Atomic and Ionic Radii absorption spectroscopy. These elements when The alkali metal atoms have the largest sizes irradiated with light, the light energy absorbed in a particular period of the periodic table. With may be sufficient to make an atom lose electron. 292 C:\ChemistryXI\Unit-10\Unit-10-Lay-3(reprint).pmd Reprint 27.7.6  THE s-BLOCK ELEMENTS 293 Table 10.1 Atomic and Physical Properties of the Alkali Metals Property Lithium Sodium Potassium Rubidium Caesium Francium Li Na K Rb Cs Fr Atomic number 3 11 19 37 55 87 –1 Atomic mass (g mol ) 6.94 22.99 39.10 85.47 132.91 (223) 1 1 1 1 1 1 Electronic [He] 2s [Ne] 3s [Ar] 4s [Kr] 5s [Xe] 6s [Rn] 7s configuration d Ionization 520 496 419 403 376 ~375 –1 enthalpy / kJ mol he Hydration –506 –406 –330 –310 –276 – –1 enthalpy/kJ mol Metallic 152 186 227 248 265 – radius / pm is Ionic radius 76 102 138 152 167 (180) + M / pm bl m.p. / K 454 371 336 312 302 – b.p / K 1615 1156 1032 961 944 – pu Density / g cm –3 Standard potentials 0 + 0.53 –3.04 0.97 –2.714 0.86 –2.925 1.53 –2.930 1.90 –2.927 – – E / V for (M / M) be T Occurrence in 18* 2.27** 1.84** 78-12* 2-6* ~ 10 –18 * re lithosphere† o R *ppm (part per million), ** percentage by weight; † Lithosphere: The Earth’s outer layer: its crust and part of the upper mantle tt E This property makes caesium and potassium 2 Na + O2 → Na 2 O 2 (peroxide) useful as electrodes in photoelectric cells. M + O 2 → MO 2 (superoxide) C 10.1.6 Chemical Properties (M = K, Rb, Cs) The alkali metals are highly reactive due to In all these oxides the oxidation state of the their large size and low ionization enthalpy. The no N alkali metal is +1. Lithium shows exceptional reactivity of these metals increases down the behaviour in reacting directly with nitrogen of group. air to form the nitride, Li3N as well. Because of (i) Reactivity towards air: The alkali metals their high reactivity towards air and water, tarnish in dry air due to the formation of © alkali metals are normally kept in kerosene oil. their oxides which in turn react with moisture to form hydroxides. They burn Problem 10.1 vigorously in oxygen forming oxides. What is the oxidation state of K in KO2? Lithium forms monoxide, sodium forms peroxide, the other metals form Solution – superoxides. The superoxide O2 ion is The superoxide species is represented as – stable only in the presence of large cations O 2 ; since the compound is neutral, such as K, Rb, Cs. therefore, the oxidation state of potassium is +1. 4 Li + O 2 → 2 Li 2 O (oxide) 293 C:\ChemistryXI\Unit-10\Unit-10-Lay-3(reprint).pmd Reprint 27.7.6  294 CHEMISTRY (ii) Reactivity towards water: The alkali the highest hydration enthalpy which 0 metals react with water to form hydroxide accounts for its high negative E value and and dihydrogen. its high reducing power. 2 M + 2H2O → 2 M+ + 2 OH− + H2 (M = an alkali metal) Problem 10.2 0 – – It may be noted that although lithium has The E for Cl 2 /Cl is +1.36, for I2/I is 0 + + most negative E value (Table 10.1), its + 0.53, for Ag /Ag is +0.79, Na /Na is + reaction with water is less vigorous than –2.71 and for Li /Li is – 3.04. Arrange that of sodium which has the least negative d 0 the following ionic species in decreasing E value among the alkali metals. This order of reducing strength: behaviour of lithium is attributed to its – – he I , Ag, Cl , Li, Na small size and very high hydration energy. Other metals of the group react explosively Solution – – with water. The order is Li > Na > I > Ag > Cl They also react with proton donors such is as alcohol, gaseous ammonia and alkynes. (vi) Solutions in liquid ammonia: The alkali (iii) Reactivity towards dihydrogen: The metals dissolve in liquid ammonia giving alkali metals react with dihydrogen at deep blue solutions which are conducting bl about 673K (lithium at 1073K) to form in nature. hydrides. All the alkali metal hydrides are M + (x + y)NH3 →[M(NH3 )x ]+ + [e(NH3 )y ]− ionic solids with high melting points. pu The blue colour of the solution is due to the ammoniated electron which absorbs 2 M + H2 → 2 M + H − energy in the visible region of light and thus (iv) Reactivity towards halogens : The alkali imparts blue colour to the solution. The be T metals readily react vigorously with solutions are paramagnetic and on + – re halogens to form ionic halides, M X. standing slowly liberate hydrogen resulting o R However, lithium halides are somewhat in the formation of amide. covalent. It is because of the high polarisation capability of lithium ion (The M + (am) + e − + NH3 (1) → MNH2(am) + ½H2 (g) tt E distortion of electron cloud of the anion by (where ‘am’ denotes solution in ammonia.) + the cation is called polarisation). The Li ion In concentrated solution, the blue colour is very small in size and has high tendency changes to bronze colour and becomes C to distort electron cloud around the diamagnetic. negative halide ion. Since anion with large size can be easily distorted, among halides, 10.1.7 Uses no N lithium iodide is the most covalent in Lithium metal is used to make useful alloys, nature. for example with lead to make ‘white metal’ (v) Reducing nature: The alkali metals are bearings for motor engines, with aluminium strong reducing agents, lithium being the © to make aircraft parts, and with magnesium most and sodium the least powerful to make armour plates. It is used in (Table 10.1). The standard electrode thermonuclear reactions. Lithium is also used 0 potential (E ) which measures the reducing to make electrochemical cells. Sodium is used power represents the overall change : to make a Na/Pb alloy needed to make PbEt4 M(s) → M(g) sublimationenthalpy and PbMe4. These organolead compounds were + M(g) → M (g) + e − ionizationenthalpy earlier used as anti-knock additives to petrol, but nowadays vehicles use lead-free petrol. M+ (g) + H2O → M+ (aq) hydrationenthalpy Liquid sodium metal is used as a coolant in With the small size of its ion, lithium has fast breeder nuclear reactors. Potassium has 294 C:\ChemistryXI\Unit-10\Unit-10-Lay-3(reprint).pmd Reprint 27.7.6  THE s-BLOCK ELEMENTS 295 a vital role in biological systems. Potassium The hydroxides which are obtained by the chloride is used as a fertilizer. Potassium reaction of the oxides with water are all white hydroxide is used in the manufacture of soft crystalline solids. The alkali metal hydroxides soap. It is also used as an excellent absorbent are the strongest of all bases and dissolve freely of carbon dioxide. Caesium is used in devising in water with evolution of much heat on photoelectric cells. account of intense hydration. 10.2 GENERAL CHARACTERISTICS OF 10.2.2 Halides THE COMPOUNDS OF THE ALKALI The alkali metal halides, MX, (X=F,Cl,Br,I) are d METALS all high melting, colourless crystalline solids. All the common compounds of the alkali metals They can be prepared by the reaction of the he are generally ionic in nature. General appropriate oxide, hydroxide or carbonate with characteristics of some of their compounds are aqueous hydrohalic acid (HX). All of these discussed here. halides have high negative enthalpies of 0 10.2.1 Oxides and Hydroxides formation; the Δf H values for fluorides is become less negative as we go down the group, On combustion in excess of air, lithium forms 0 whilst the reverse is true for Δf H for chlorides, mainly the oxide, Li2O (plus some peroxide bromides and iodides. For a given metal Li2O2), sodium forms the peroxide, Na2O2 (and bl 0 Δf H always becomes less negative from some superoxide NaO2) whilst potassium, fluoride to iodide. rubidium and caesium form the superoxides, MO 2. Under appropriate conditions pure pu The melting and boiling points always compounds M2O, M 2O 2 and MO 2 may be follow the trend: fluoride > chloride > bromide prepared. The increasing stability of the > iodide. All these halides are soluble in water. The low solubility of LiF in water is due to its be T peroxide or superoxide, as the size of the metal ion increases, is due to the stabilisation of large high lattice enthalpy whereas the low solubility re anions by larger cations through lattice energy of CsI is due to smaller hydration enthalpy of o R effects. These oxides are easily hydrolysed by its two ions. Other halides of lithium are soluble water to form the hydroxides according to the in ethanol, acetone and ethylacetate; LiCl is following reactions : soluble in pyridine also. tt E M2O + H2O → 2M+ + 2 OH – 10.2.3 Salts of Oxo-Acids Oxo-acids are those in which the acidic proton C M2O2 + 2H2O → 2M + + 2 OH – + H2O2 is on a hydroxyl group with an oxo group attached to the same atom e.g., carbonic acid, 2 MO2 + 2 H2O → 2M+ + 2 OH – + H2O2 + O2 no N H 2 CO 3 (OC(OH) 2 ; sulphuric acid, H 2 SO 4 The oxides and the peroxides are colourless (O2S(OH)2). The alkali metals form salts with when pure, but the superoxides are yellow or all the oxo-acids. They are generally soluble orange in colour. The superoxides are also in water and thermally stable. Their © paramagnetic. Sodium peroxide is widely used carbonates (M2CO3) and in most cases the as an oxidising agent in inorganic chemistry. hydrogencarbonates (MHCO3) also are highly stable to heat. As the electropositive character Problem 10.3 increases down the group, the stability of the Why is KO2 paramagnetic ? carbonates and hydorgencarbonates increases. Solution Lithium carbonate is not so stable to heat; – lithium being very small in size polarises a The superoxide O 2 is paramagnetic 2– large CO3 ion leading to the formation of more because of one unpaired electron in π*2p stable Li2O and CO2. Its hydrogencarbonate molecular orbital. does not exist as a solid. 295 C:\ChemistryXI\Unit-10\Unit-10-Lay-3(reprint).pmd Reprint 27.7.6  296 CHEMISTRY 10.3 ANOMALOUS PROPERTIES OF and lighter than other elements in the LITHIUM respective groups. The anomalous behaviour of lithium is due to (ii) Lithium and magnesium react slowly with the : (i) exceptionally small size of its atom and water. Their oxides and hydroxides are ion, and (ii) high polarising power (i.e., charge/ much less soluble and their hydroxides radius ratio). As a result, there is increased decompose on heating. Both form a nitride, covalent character of lithium compounds which Li3N and Mg3N2, by direct combination is responsible for their solubility in organic with nitrogen. solvents. Further, lithium shows diagonal (iii) The oxides, Li2O and MgO do not combine d relationship to magnesium which has been with excess oxygen to give any superoxide. discussed subsequently. (iv) The carbonates of lithium and magnesium he 10.3.1 Points of Difference between decompose easily on heating to Lithium and other Alkali Metals form the oxides and CO 2. Solid hydrogencarbonates are not formed by (i) Lithium is much harder. Its m.p. and b.p. lithium and magnesium. are higher than the other alkali metals. is (v) Both LiCl and MgCl2 are soluble in ethanol. (ii) Lithium is least reactive but the strongest reducing agent among all the alkali metals. (vi) Both LiCl and MgCl2 are deliquescent and crystallise from aqueous solution as bl On combustion in air it forms mainly monoxide, Li2O and the nitride, Li3N unlike hydrates, LiCl·2H2O and MgCl2·8H2O. other alkali metals. pu 10.4 SOME IMPORTANT COMPOUNDS OF (iii) LiCl is deliquescent and crystallises as a SODIUM hydrate, LiCl.2H2O whereas other alkali Industrially important compounds of sodium metal chlorides do not form hydrates. include sodium carbonate, sodium hydroxide, be T (iv) Lithium hydrogencarbonate is not sodium chloride and sodium bicarbonate. The obtained in the solid form while all other large scale production of these compounds re o R elements form solid hydrogencarbonates. and their uses are described below: (v) Lithium unlike other alkali metals forms Sodium Carbonate (Washing Soda), no ethynide on reaction with ethyne. Na2CO3·10H2O tt E (vi) Lithium nitrate when heated gives lithium Sodium carbonate is generally prepared by oxide, Li2O, whereas other alkali metal Solvay Process. In this process, advantage is C nitrates decompose to give the taken of the low solubility of sodium corresponding nitrite. hydrogencarbonate whereby it gets precipitated in the reaction of sodium chloride 4LiNO 3 → 2 Li 2 O + 4 NO 2 + O 2 no N with ammonium hydrogencarbonate. The 2 NaNO 3 → 2 NaNO 2 + O 2 latter is prepared by passing CO 2 to a concentrated solution of sodium chloride (vii) LiF and Li2O are comparatively much less saturated with ammonia, where ammonium soluble in water than the corresponding © carbonate followed by ammonium compounds of other alkali metals. hydrogencarbonate are formed. The equations 10.3.2 Points of Similarities between for the complete process may be written as : Lithium and Magnesium 2 NH 3 + H 2 O + CO 2 → ( NH 4 )2 CO 3 The similarity between lithium and magnesium ( NH 4 )2 CO 3 + H 2O + CO2 → 2 NH 4 HCO 3 is particularly striking and arises because of their similar sizes : atomic radii, Li = 152 pm, NH 4 HCO 3 + NaCl → NH 4 Cl + NaHCO 3 + Mg = 160 pm; ionic radii : Li = 76 pm, Sodium hydrogencarbonate crystal 2+ Mg = 72 pm. The main points of similarity are: separates. These are heated to give sodium (i) Both lithium and magnesium are harder carbonate. 296 C:\ChemistryXI\Unit-10\Unit-10-Lay-3(reprint).pmd Reprint 27.7.6  THE s-BLOCK ELEMENTS 297 2 NaHCO 3 → Na 2 CO 3 + CO 2 + H 2 O of brine solution, contains sodium sulphate, calcium sulphate, calcium chloride and In this process NH3 is recovered when the magnesium chloride as impurities. Calcium solution containing NH4Cl is treated with chloride, CaCl2, and magnesium chloride, Ca(OH)2. Calcium chloride is obtained as a MgCl 2 are impurities because they are by-product. deliquescent (absorb moisture easily from the 2 NH 4 Cl + Ca ( OH )2 → 2 NH3 + CaCl 2 + H2 O atmosphere). To obtain pure sodium chloride, It may be mentioned here that Solvay the crude salt is dissolved in minimum amount of water and filtered to remove insoluble d process cannot be extended to the manufacture of potassium carbonate because impurities. The solution is then saturated with potassium hydrogencarbonate is too soluble hydrogen chloride gas. Crystals of pure he to be precipitated by the addition of sodium chloride separate out. Calcium and ammonium hydrogencarbonate to a saturated magnesium chloride, being more soluble than solution of potassium chloride. sodium chloride, remain in solution. Properties : Sodium carbonate is a white Sodium chloride melts at 1081K. It has a is crystalline solid which exists as a decahydrate, solubility of 36.0 g in 100 g of water at 273 K. Na2CO3·10H2O. This is also called washing The solubility does not increase appreciably soda. It is readily soluble in water. On heating, with increase in temperature. bl the decahydrate loses its water of crystallisation to form monohydrate. Above 373K, the Uses : monohydrate becomes completely anhydrous pu (i) It is used as a common salt or table salt for and changes to a white powder called soda ash. domestic purpose. Na2CO3 10H2O ⎯⎯⎯⎯ 375 K → Na2CO3 H2O + 9H2O (ii) It is used for the preparation of Na2O2, NaOH and Na2CO3. be T >373K Na 2CO3 H2O ⎯⎯⎯⎯ → Na 2CO3 + H2O Sodium Hydroxide (Caustic Soda), NaOH re Carbonate part of sodium carbonate gets o R hydrolysed by water to form an alkaline Sodium hydroxide is generally prepared solution. commercially by the electrolysis of sodium chloride in Castner-Kellner cell. A brine 3 + H2 O → HCO3 + OH CO2– – – tt E solution is electrolysed using a mercury Uses: cathode and a carbon anode. Sodium metal (i) It is used in water softening, laundering discharged at the cathode combines with C and cleaning. mercury to form sodium amalgam. Chlorine (ii) It is used in the manufacture of glass, gas is evolved at the anode. soap, borax and caustic soda. no N (iii) It is used in paper, paints and textile Cathode : Na + + e − ⎯⎯⎯Hg → Na – amalgam industries. 1 Anode : Cl – → Cl 2 + e – (iv) It is an important laboratory reagent both 2 © in qualitative and quantitative analysis. The amalgam is treated with water to give sodium hydroxide and hydrogen gas. Sodium Chloride, NaCl 2Na-amalgam + 2H2OÆ2NaOH+ 2Hg +H2 The most abundant source of sodium chloride is sea water which contains 2.7 to 2.9% by Sodium hydroxide is a white, translucent mass of the salt. In tropical countries like India, solid. It melts at 591 K. It is readily soluble in common salt is generally obtained by water to give a strong alkaline solution. evaporation of sea water. Approximately 50 Crystals of sodium hydroxide are deliquescent. lakh tons of salt are produced annually in The sodium hydroxide solution at the surface India by solar evaporation. Crude sodium reacts with the CO2 in the atmosphere to form chloride, generally obtained by crystallisation Na2CO3. 297 C:\ChemistryXI\Unit-10\Unit-10-Lay-3(reprint).pmd Reprint 27.7.6  298 CHEMISTRY Uses: It is used in (i) the manufacture of soap, found on the opposite sides of cell membranes. paper, artificial silk and a number of chemicals, As a typical example, in blood plasma, sodium –1 (ii) in petroleum refining, (iii) in the purification is present to the extent of 143 mmolL , of bauxite, (iv) in the textile industries for whereas the potassium level is only –1 mercerising cotton fabrics, (v) for the 5 mmolL within the red blood cells. These –1 + preparation of pure fats and oils, and (vi) as a concentrations change to 10 mmolL (Na ) and –1 + laboratory reagent. 105 mmolL (K ). These ionic gradients Sodium Hydrogencarbonate (Baking demonstrate that a discriminatory mechanism, Soda), NaHCO3 called the sodium-potassium pump, operates d across the cell membranes which consumes Sodium hydrogencarbonate is known as more than one-third of the ATP used by a baking soda because it decomposes on heating he resting animal and about 15 kg per 24 h in a to generate bubbles of carbon dioxide (leaving resting human. holes in cakes or pastries and making them light and fluffy). 10.6 GROUP 2 ELEMENTS : ALKALINE Sodium hydrogencarbonate is made by EARTH METALS is saturating a solution of sodium carbonate with The group 2 elements comprise beryllium, carbon dioxide. The white crystalline powder magnesium, calcium, strontium, barium and of sodium hydrogencarbonate, being less radium. They follow alkali metals in the bl soluble, gets separated out. periodic table. These (except beryllium) are known as alkaline earth metals. The first Na 2 CO3 H2O CO 2 2 NaHCO 3 pu element beryllium differs from the rest of the Sodium hydrogencarbonate is a mild members and shows diagonal relationship to antiseptic for skin infections. It is used in fire aluminium. The atomic and physical extinguishers. properties of the alkaline earth metals are be T shown in Table 10.2. 10.5 BIOLOGICAL IMPORTANCE OF re SODIUM AND POTASSIUM 10.6.1 Electronic Configuration o R A typical 70 kg man contains about 90 g of Na These elements have two electrons in the and 170 g of K compared with only 5 g of iron s -orbital of the valence shell (Table 10.2). Their tt E and 0.06 g of copper. general electronic configuration may be 2 Sodium ions are found primarily on the represented as [noble gas] ns. Like alkali outside of cells, being located in blood plasma metals, the compounds of these elements are C and in the interstitial fluid which surrounds also predominantly ionic. the cells. These ions participate in the Element Symbol Electronic transmission of nerve signals, in regulating the configuration no N flow of water across cell membranes and in the transport of sugars and amino acids into cells. Beryllium Be 1s22s2 Sodium and potassium, although so similar Magnesium Mg 1s22s22p63s2 chemically, differ quantitatively in their ability Calcium Ca 1s22s22p63s23p64s2 © to penetrate cell membranes, in their transport Strontium Sr 1s22s22p63s23p63d10 mechanisms and in their efficiency to activate 4s24p65s2 enzymes. Thus, potassium ions are the most Barium Ba 1s22s22p63s23p63d104s2 abundant cations within cell fluids, where they 4p64d105s25p66s2 or activate many enzymes, participate in the [Xe]6s2 oxidation of glucose to produce ATP and, with Radium Ra [Rn]7s2 sodium, are responsible for the transmission of nerve signals. 10.6.2 Atomic and Ionic Radii There is a very considerable variation in the The atomic and ionic radii of the alkaline earth concentration of sodium and potassium ions metals are smaller than those of the 298 C:\ChemistryXI\Unit-10\Unit-10-Lay-3(reprint).pmd Reprint 27.7.6  THE s-BLOCK ELEMENTS 299 Table 10.2 Atomic and Physical Properties of the Alkaline Earth Metals Property Beryllium Magnesium Calcium Strontium Barium Radium Be Mg Ca Sr Ba Ra Atomic number 4 12 20 38 56 88 –1 Atomic mass (g mol ) 9.01 24.31 40.08 87.62 137.33 226.03 Electronic [He] 2s2 [Ne] 3s2 [Ar] 4s2 [Kr] 5s2 [Xe] 6s2 [Rn] 7s2 configuration d Ionization 899 737 590 549 503 509 enthalpy (I) / kJ mol–1 Ionization 1757 1450 1145 1064 965 979 he enthalpy (II) /kJ mol–1 Hydration enthalpy – 2494 – 1921 –1577 – 1443 – 1305 – (kJ/mol) Metallic 111 160 197 215 222 – is radius / pm Ionic radius 31 72 100 118 135 148 M2+ / pm bl m.p. / K 1560 924 1124 1062 1002 973 b.p / K 2745 1363 1767 1655 2078 (1973) Density / g cm–3 pu 1.84 1.74 1.55 2.63 3.59 (5.5) Standard potential –1.97 –2.36 –2.84 –2.89 – 2.92 –2.92 E0 / V for (M2+/ M) be T Occurrence in 2* 2.76** 4.6** 384* 390 * 10–6* lithosphere re *ppm (part per million); ** percentage by weight o R corresponding alkali metals in the same increase in ionic size down the group. periods. This is due to the increased nuclear 2+ 2+ 2+ Be > Mg > Ca > Sr > Ba 2+ 2+ tt E charge in these elements. Within the group, the The hydration enthalpies of alkaline earth atomic and ionic radii increase with increase metal ions are larger than those of alkali metal in atomic number. C ions. Thus, compounds of alkaline earth metals 10.6.3 Ionization Enthalpies are more extensively hydrated than those of The alkaline earth metals have low ionization alkali metals, e.g., MgCl2 and CaCl2 exist as no N enthalpies due to fairly large size of the atoms. MgCl2.6H2O and CaCl2· 6H2O while NaCl and Since the atomic size increases down the KCl do not form such hydrates. group, their ionization enthalpy decreases 10.6.5 Physical Properties (Table 10.2). The first ionisation enthalpies of The alkaline earth metals, in general, are silvery © the alkaline earth metals are higher than those white, lustrous and relatively soft but harder of the corresponding Group 1 metals. This is than the alkali metals. Beryllium and due to their small size as compared to the magnesium appear to be somewhat greyish. corresponding alkali metals. It is interesting The melting and boiling points of these metals to note that the second ionisation enthalpies are higher than the corresponding alkali metals of the alkaline earth metals are smaller than due to smaller sizes. The trend is, however, not those of the corresponding alkali metals. systematic. Because of the low ionisation 10.6.4 Hydration Enthalpies enthalpies, they are strongly electropositive in Like alkali metal ions, the hydration enthalpies nature. The electropositive character increases of alkaline earth metal ions decrease with down the group from Be to Ba. Calcium, 299 299 C:\ChemistryXI\Unit-10\Unit-10-Lay-3(reprint).pmd 27.7.6, C:\ChemistryXI\Unit-10\Unit-10-Lay-3(reprint).pmd Reprint16.10.6 27.7.6(reprint)  300 CHEMISTRY strontium and barium impart characteristic (iv) Reactivity towards acids: The alkaline brick red, crimson and apple green colours earth metals readily react with acids liberating respectively to the flame. In flame the electrons dihydrogen. are excited to higher energy levels and when M + 2HCl → MCl2 + H2 they drop back to the ground state, energy is (v) Reducing nature: Like alkali metals, the emitted in the form of visible light. The alkaline earth metals are strong reducing electrons in beryllium and magnesium are too strongly bound to get excited by flame. Hence, agents. This is indicated by large negative these elements do not impart any colour to the values of their reduction potentials d flame. The flame test for Ca, Sr and Ba is (Table 10.2). However their reducing power is helpful in their detection in qualitative analysis less than those of their corresponding alkali metals. Beryllium has less negative value he and estimation by flame photometry. The alkaline earth metals like those of alkali metals compared to other alkaline earth metals. have high electrical and thermal conductivities However, its reducing nature is due to large which are typical characteristics of metals. hydration energy associated with the small 2+ size of Be ion and relatively large value of the 10.6.6 Chemical Properties is atomization enthalpy of the metal. The alkaline earth metals are less reactive than (vi) Solutions in liquid ammonia: Like the alkali metals. The reactivity of these alkali metals, the alkaline earth metals dissolve bl elements increases on going down the group. in liquid ammonia to give deep blue black (i) Reactivity towards air and water: solutions forming ammoniated ions. Beryllium and magnesium are kinetically inert pu 2+ M + ( x + y ) NH3 → ⎡⎣M ( NH3 ) X ⎤⎦ + 2 ⎡⎣e ( NH3 ) Y ⎤⎦ – to oxygen and water because of the formation of an oxide film on their surface. However, powdered beryllium burns brilliantly on From these solutions, the ammoniates, be T 2+ ignition in air to give BeO and Be 3 N 2. [M(NH3)6] can be recovered. Magnesium is more electropositive and burns re with dazzling brilliance in air to give MgO and 10.6.7 Uses o R Mg3N2. Calcium, strontium and barium are Beryllium is used in the manufacture of alloys. readily attacked by air to form the oxide and Copper-beryllium alloys are used in the nitride. They also react with water with tt E preparation of high strength springs. Metallic increasing vigour even in cold to form beryllium is used for making windows of hydroxides. X-ray tubes. Magnesium forms alloys with C (ii) Reactivity towards the halogens: All aluminium, zinc, manganese and tin. the alkaline earth metals combine with halogen Magnesium-aluminium alloys being light in at elevated temperatures forming their halides. mass are used in air-craft construction. no N M + X 2 → MX 2 ( X = F, Cl, Br, l ) Magnesium (powder and ribbon) is used in Thermal decomposition of (NH4)2BeF4 is the flash powders and bulbs, incendiary bombs best route for the preparation of BeF2, and and signals. A suspension of magnesium BeCl2 is conveniently made from the oxide. hydroxide in water (called milk of magnesia) © 600 − 800K is used as antacid in medicine. Magnesium BeO + C + Cl 2 BeCl 2 + CO carbonate is an ingredient of toothpaste. (iii) Reactivity towards hydrogen: All the Calcium is used in the extraction of metals from elements except beryllium combine with oxides which are difficult to reduce with hydrogen upon heating to form their hydrides, carbon. Calcium and barium metals, owing MH2. to their reactivity with oxygen and nitrogen at BeH2, however, can be prepared by the reaction elevated temperatures, have often been used of BeCl2 with LiAlH4. to remove air from vacuum tubes. Radium salts are used in radiotherapy, for example, in 2BeCl 2 + LiAlH 4 → 2BeH 2 + LiCl + AlCl 3 the treatment of cancer. 300 C:\ChemistryXI\Unit-10\Unit-10-Lay-3(reprint).pmd Reprint 27.7.6  THE s-BLOCK ELEMENTS 301 10.7 GENERAL CHARACTERISTICS OF In the vapour phase BeCl2 tends to form a COMPOUNDS OF THE ALKALINE chloro-bridged dimer which dissociates into the EARTH METALS linear monomer at high temperatures of the 2+ The dipositive oxidation state (M ) is the order of 1200 K. The tendency to form halide predominant valence of Group 2 elements. The hydrates gradually decreases (for example, alkaline earth metals form compounds which MgCl2·8H2O, CaCl2·6H2O, SrCl2·6H2O and are predominantly ionic but less ionic than the BaCl2·2H2O) down the group. The dehydration corresponding compounds of alkali metals. of hydrated chlorides, bromides and iodides This is due to increased nuclear charge and of Ca, Sr and Ba can be achieved on heating; d smaller size. The oxides and other compounds however, the corresponding hydrated halides of beryllium and magnesium are more covalent of Be and Mg on heating suffer hydrolysis. The he than those formed by the heavier and large fluorides are relatively less soluble than the sized members (Ca, Sr, Ba). The general chlorides owing to their high lattice energies. characteristics of some of the compounds of (iii) Salts of Oxoacids: The alkaline earth alkali earth metals are described below. metals also form salts of oxoacids. Some of is (i) Oxides and Hydroxides: The alkaline these are : earth metals burn in oxygen to form the Carbonates: Carbonates of alkaline earth monoxide, MO which, except for BeO, have metals are insoluble in water and can be bl rock-salt structure. The BeO is essentially precipitated by addition of a sodium or covalent in nature. The enthalpies of formation ammonium carbonate solution to a solution of these oxides are quite high and consequently of a soluble salt of these metals. The solubility pu they are very stable to heat. BeO is amphoteric while oxides of other elements are ionic in of carbonates in water decreases as the atomic number of the metal ion increases. All the nature. All these oxides except BeO are basic be T in nature and react with water to form sparingly carbonates decompose on heating to give soluble hydroxides. carbon dioxide and the oxide. Beryllium re carbonate is unstable and can be kept only in o R MO + H2O → M(OH)2 the atmosphere of CO2. The thermal stability The solubility, thermal stability and the increases with increasing cationic size. basic character of these hydroxides increase tt E Sulphates: The sulphates of the alkaline earth with increasing atomic number from Mg(OH)2 to Ba(OH) 2. The alkaline earth metal metals are all white solids and stable to heat. hydroxides are, however, less basic and less BeSO4, and MgSO4 are readily soluble in water; C stable than alkali metal hydroxides. Beryllium the solubility decreases from CaSO4 to BaSO4. 2+ hydroxide is amphoteric in nature as it reacts The greater hydration enthalpies of Be and 2+ with acid and alkali both. Mg ions overcome the lattice enthalpy factor no N – 2– and therefore their sulphates are soluble in Be(OH)2 + 2OH → [Be(OH)4] Beryllate ion water. Be(OH)2 + 2HCl + 2H2O → [Be(OH)4]Cl2 Nitrates: The nitrates are made by dissolution © of the carbonates in dilute nitric acid. (ii) Halides: Except for beryllium halides, all Magnesium nitrate crystallises with six other halides of alkaline earth metals are ionic molecules of water, whereas barium nitrate in nature. Beryllium halides are essentially crystallises as the anhydrous salt. This again covalent and soluble in organic solvents. shows a decreasing tendency to form hydrates Beryllium chloride has a chain structure in the with increasing size and decreasing hydration solid state as shown below: enthalpy. All of them decompose on heating to give the oxide like lithium nitrate. 2M ( NO 3 )2 → 2MO + 4NO2 + O2 (M = Be, Mg, Ca, Sr, Ba) 301 C:\ChemistryXI\Unit-10\Unit-10-Lay-3(reprint).pmd Reprint 27.7.6  302 CHEMISTRY Problem 10.4 (iii) The oxide and hydroxide of beryllium, unlike the hydroxides of other elements in Why does the solubility of alkaline earth the group, are amphoteric in nature. metal hydroxides in water increase down the group? 10.8.1 Diagonal Relationship between Beryllium and Aluminium Solution 2+ Among alkaline earth metal hydroxides, The ionic radius of Be is estimated to be the anion being common the cationic 31 pm; the charge/radius ratio is nearly the 3+ radius will influence the lattice enthalpy. same as that of the Al ion. Hence beryllium d Since lattice enthalpy decreases much resembles aluminium in some ways. Some of more than the hydration enthalpy with the similarities are: he increasing ionic size, the solubility (i) Like aluminium, beryllium is not readily increases as we go down the group. attacked by acids because of the presence of an oxide film on the surface of the metal. Problem 10.5 (ii) Beryllium hydroxide dissolves in excess of Why does the solubility of alkaline earth 2– is alkali to give a beryllate ion, [Be(OH)4] just metal carbonates and sulphates in water as aluminium hydroxide gives aluminate decrease down the group? ion, [Al(OH)4]. – bl Solution (iii) The chlorides of both beryllium and – The size of anions being much larger aluminium have Cl bridged chloride compared to cations, the lattice enthalpy pu structure in vapour phase. Both the will remain almost constant within a chlorides are soluble in organic solvents particular group. Since the hydration and are strong Lewis acids. They are used enthalpies decrease down the group, as Friedel Craft catalysts. be T solubility will decrease as found for (iv) Beryllium and aluminium ions have strong alkaline earth metal carbonates and 2– tendency to form complexes, BeF4 , AlF6. 3– re o R sulphates. 10.9 SOME IMPORTANT COMPOUNDS OF CALCIUM 10.8 ANOMALOUS BEHAVIOUR OF tt E Important compounds of calcium are calcium BERYLLIUM oxide, calcium hydroxide, calcium sulphate, Beryllium, the first member of the Group 2 calcium carbonate and cement. These are C metals, shows anomalous behaviour as industrially important compounds. The large compared to magnesium and rest of the scale preparation of these compounds and members. Further, it shows diagonal their uses are described below. no N relationship to aluminium which is discussed Calcium Oxide or Quick Lime, CaO subsequently. It is prepared on a commercial scale by (i) Beryllium has exceptionally small atomic heating limestone (CaCO3) in a rotary kiln at and ionic sizes and thus does not compare © well with other members of the group. 1070-1270 K. Because of high ionisation enthalpy and CaCO3 heat CaO + CO2 small size it forms compounds which are largely covalent and get easily hydrolysed. The carbon dioxide is removed as soon as (ii) Beryllium does not exhibit coordination it is produced to enable the reaction to proceed number more than four as in its valence to completion. shell there are only four orbitals. The Calcium oxide is a white amorphous solid. remaining members of the group can have It has a melting point of 2870 K. On exposure a coordination number of six by making to atmosphere, it absorbs moisture and carbon use of d-orbitals. dioxide. 302 C:\ChemistryXI\Unit-10\Unit-10-Lay-3(reprint).pmd Reprint 27.7.6  THE s-BLOCK ELEMENTS 303 CaO + H2 O → Ca ( OH )2 (ii) It is used in white wash due to its disinfectant nature. CaO + CO 2 → CaCO 3 (iii) It is used in glass making, in tanning The addition of limited amount of water industry, for the preparation of bleaching breaks the lump of lime. This process is called powder and for purification of sugar. slaking of lime. Quick lime slaked with soda Calcium Carbonate, CaCO3 gives solid sodalime. Being a basic oxide, it Calcium carbonate occurs in nature in several combines with acidic oxides at high forms like limestone, chalk, marble etc. It can temperature. d be prepared by passing carbon dioxide CaO + SiO 2 → CaSiO 3 through slaked lime or by the addition of sodium carbonate to calcium chloride. he 6CaO + P4 O10 → 2Ca 3 ( PO 4 )2 Ca ( OH )2 + CO2 → CaCO3 + H2 O Uses: (i) It is an important primary material for CaCl 2 + Na 2 CO 3 → CaCO 3 + 2NaCl manufacturing cement and is the cheapest Excess of carbon dioxide should be is form of alkali. avoided since this leads to the formation of (ii) It is used in the manufacture of sodium water soluble calcium hydrogencarbonate. Calcium carbonate is a white fluffy powder. bl carbonate from caustic soda. It is almost insoluble in water. When heated (iii) It is employed in the purification of sugar to 1200 K, it decomposes to evolve carbon and in the manufacture of dye stuffs. pu dioxide. Calcium Hydroxide (Slaked lime), Ca(OH)2 CaCO3 ⎯⎯⎯⎯ 1200 K → CaO + CO2 Calcium hydroxide is prepared by adding It reacts with dilute acid to liberate carbon be T water to quick lime, CaO. dioxide. It is a white amorphous powder. It is re CaCO 3 + 2HCl → CaCl 2 + H 2 O + CO 2 o R sparingly soluble in water. The aqueous solution is known as lime water and a CaCO 3 + H 2 SO 4 → CaSO 4 + H 2 O + CO 2 suspension of slaked lime in water is known Uses: tt E as milk of lime. It is used as a building material in the form of When carbon dioxide is passed through marble and in the manufacture of quick lime. lime water it turns milky due to the formation Calcium carbonate along with magnesium C of calcium carbonate. carbonate is used as a flux in the extraction of metals such as iron. Specially precipitated Ca ( OH )2 + CO2 → CaCO 3 + H2 O CaCO3 is extensively used in the manufacture no N On passing excess of carbon dioxide, the of high quality paper. It is also used as an precipitate dissolves to form calcium antacid, mild abrasive in tooth paste, a hydrogencarbonate. constituent of chewing gum, and a filler in cosmetics. © CaCO3 + CO2 + H2 O → Ca ( HCO3 )2 Calcium Sulphate (Plaster of Paris), Milk of lime reacts with chlorine to form CaSO4· H2O hypochlorite, a constituent of bleaching It is a hemihydrate of calcium sulphate. It is powder. obtained when gypsum, CaSO 4·2H 2 O, is heated to 393 K. 2Ca ( OH )2 + 2Cl2 → CaCl 2 + Ca OCl ( ) 2 + 2H2O 2 ( CaSO4.2H2O ) → 2 ( CaSO4 ).H2O + 3H2 O Bleaching powder Uses: Above 393 K, no water of crystallisation is left (i) It is used in the preparation of mortar, a and anhydrous calcium sulphate, CaSO4 is building material. formed. This is known as ‘dead burnt plaster’. 303 C:\ChemistryXI\Unit-10\Unit-10-Lay-3(reprint).pmd Reprint 27.7.6  304 CHEMISTRY It has a remarkable property of setting with silicate (Ca 3 SiO 5 ) 51% and tricalcium water. On mixing with an adequate quantity aluminate (Ca3Al2O6) 11%. of water it forms a plastic mass that gets into a Setting of Cement: When mixed with water, hard solid in 5 to 15 minutes. the setting of cement takes place to give a hard Uses: mass. This is due to the hydration of the The largest use of Plaster of Paris is in the molecules of the constituents and their building industry as well as plasters. It is used rearrangement. The purpose of adding for immoblising the affected part of organ where gypsum is only to slow down the process of there is a bone fracture or sprain. It is also setting of the cement so that it gets sufficiently d employed in dentistry, in ornamental work and hardened. for making casts of statues and busts. Uses: Cement has become a commodity of he Cement: Cement is an important building national necessity for any country next to iron material. It was first introduced in England in and steel. It is used in concrete and reinforced 1824 by Joseph Aspdin. It is also called concrete, in plastering and in the construction Portland cement because it resembles with the of bridges, dams and buildings. is natural limestone quarried in the Isle of 10.10 BIOLOGICAL IMPORTANCE OF Portland, England. MAGNESIUM AND CALCIUM Cement is a product obtained by bl An adult body contains about 25 g of Mg and combining a material rich in lime, CaO with 1200 g of Ca compared with only 5 g of iron other material such as clay which contains and 0.06 g of copper. The daily requirement silica, SiO 2 along with the oxides of pu in the human body has been estimated to be aluminium, iron and magnesium. The average 200 – 300 mg. composition of Portland cement is : CaO, 50- All enzymes that utilise ATP in phosphate 60%; SiO2, 20-25%; Al2O3, 5-10%; MgO, 2- be T transfer require magnesium as the cofactor. 3%; Fe2O3, 1-2% and SO3, 1-2%. For a good The main pigment for the absorption of light re quality cement, the ratio of silica (SiO2) to in plants is chlorophyll which contains o R alumina (Al2O3) should be between 2.5 and 4 magnesium. About 99 % of body calcium is and the ratio of lime (CaO) to the total of the present in bones and teeth. It also plays oxides of silicon (SiO2) aluminium (Al2O3) important roles in neuromuscular function, tt E and iron (Fe2O3) should be as close as possible interneuronal transmission, cell membrane to 2. integrity and blood coagulation. The calcium The raw materials for the manufacture of C concentration in plasma is regulated at about cement are limestone and clay. When clay and 100 mgL–1. It is maintained by two hormones: lime are strongly heated together they fuse and calcitonin and parathyroid hormone. Do you no N react to form ‘cement clinker’. This clinker is know that bone is not an inert and unchanging mixed with 2-3% by weight of gypsum substance but is continuously being (CaSO4·2H2O) to form cement. Thus important solubilised and redeposited to the extent of ingredients present in Portland cement are 400 mg per day in man? All this calcium dicalcium silicate (Ca2SiO4) 26%, tricalcium © passes through the plasma. SUMMARY The s-Block of the periodic table constitutes Group1 (alkali metals) and Group 2 (alkaline earth metals). They are so called because their oxides and hydroxides are alkaline in nature. The alkali metals are characterised by one s-electron and the alkaline earth metals by two s-electrons in the+ valence shell of their atoms. These are highly reactive 2+ metals forming monopositive (M ) and dipositve (M ) ions respectively. 304 C:\ChemistryXI\Unit-10\Unit-10-Lay-3(reprint).pmd Reprint 27.7.6  THE s-BLOCK ELEMENTS 305 There is a regular trend in the physical and chemical properties of the alkali metal with increasing atomic numbers. The atomic and ionic sizes increase and the ionization enthalpies decrease systematically down the group. Somewhat similar trends are observed among the properties of the alkaline earth metals. The first element in each of these groups, lithium in Group 1 and beryllium in Group 2 shows similarities in properties to the second member of the next group. Such similarities are termed as the ‘diagonal relationship’ in the periodic table. As such these elements are anomalous as far as their group characteristics are concerned. The alkali metals are silvery white, soft and low melting. They are highly reactive. d The compounds of alkali metals are predominantly ionic. Their oxides and hydroxides are soluble in water forming strong alkalies. Important compounds of sodium includes sodium carbonate, sodium chloride, sodium hydroxide and sodium hydrogencarbonate. he Sodium hydroxide is manufactured by Castner-Kellner process and sodium carbonate by Solvay process. The chemistry of alkaline earth metals is very much like that of the alkali metals. However, some differences arise because of reduced atomic and ionic sizes and increased cationic charges in case of alkaline earth metals. Their oxides and hydroxides are less is basic than the alkali metal oxides and hydroxides. Industrially important compounds of calcium include calcium oxide (lime), calcium hydroxide (slaked lime), calcium sulphate (Plaster of Paris), calcium carbonate (limestone) and cement. Portland cement is an bl important constructional material. It is manufactured by heating a pulverised mixture of limestone and clay in a rotary kiln. The clinker thus obtained is mixed with some gypsum (2-3%) to give a fine powder of cement. All these substances find variety of uses pu in different areas. Monovalent sodium and potassium ions and divalent magnesium and calcium ions are found in large proportions in biological fluids. These ions perform important be T biological functions such as maintenance of ion balance and nerve impulse conduction. re o R EXERCISES tt E 10.1 What are the common physical and chemical features of alkali metals ? 10.2 Discuss the general characteristics and gradation in properties of alkaline earth metals. C 10.3 Why are alkali metals not found in nature ? 10.4 Find out the oxidation state of sodium in Na2O2. 10.5 Explain why is sodium less reactive than potassium. no N 10.6 Compare the alkali metals and alkaline earth metals with respect to (i) ionisation enthalpy (ii) basicity of oxides and (iii) solubility of hydroxides. 10.7 In what ways lithium shows similarities to magnesium in its chemical behaviour? 10.8 Explain why can alkali and alkaline earth metals not be obtained by chemical © reduction methods? 10.9 Why are potassium and caesium, rather than lithium used in photoelectric cells? 10.10 When an alkali metal dissolves in liquid ammonia the solution can acquire different colours. Explain the reasons for this type of colour change. 10.11 Beryllium and magnesium do not give colour to flame whereas other alkaline earth metals do so. Why ? 10.12 Discuss the various reactions that occur in the Solvay process. 10.13 Potassium carbonate cannot be prepared by Solvay process. Why ? 10.14 Why is Li2CO3 decomposed at a lower temperature whereas Na2CO3 at higher temperature? 305 C:\ChemistryXI\Unit-10\Unit-10-Lay-3(reprint).pmd Reprint 27.7.6  306 CHEMISTRY 10.15 Compare the solubility and thermal stability of the following compounds of the alkali metals with those of the alkaline earth metals. (a) Nitrates (b) Carbonates (c) Sulphates. 10.16 Starting with sodium chloride how would you proceed to prepare (i) sodium metal (ii) sodium hydroxide (iii) sodium peroxide (iv) sodium carbonate ? 10.17 What happens when (i) magnesium is burnt in air (ii) quick lime is heated with silica (iii) chlorine reacts with slaked lime (iv) calcium nitrate is heated ? 10.18 Describe two important uses of each of the following : (i) caustic soda (ii) sodium carbonate (iii) quicklime. d 10.19 Draw the structure of (i) BeCl2 (vapour) (ii) BeCl2 (solid). 10.20 The hydroxides and carbonates of sodium and potassium are easily soluble in he water while the corresponding salts of magnesium and calcium are sparingly soluble in water. Explain. 10.21 Describe the importance of the following : (i) limestone (ii) cement (iii) plaster of paris. 10.22 Why are lithium salts commonly hydrated and those of the other alkali ions is usually anhydrous? 10.23 Why is LiF almost insoluble in water whereas LiCl soluble not only in water but also in acetone ? bl 10.24 Explain the significance of sodium, potassium, magnesium and calcium in biological fluids. 10.25 pu What happens when (i) sodium metal is dropped in water ? (ii) sodium metal is heated in free supply of air ? (iii) sodium peroxide dissolves in water ? be T 10.26 Comment on each of the following observations: re + + + (a) The mobilities of the alkali metal ions in aqueous solution are Li < Na < K o R + + < Rb < Cs (b) Lithium is the only alkali metal to form a nitride directly. 0 2+ – (c) E for M (aq) + 2e → M(s) (where M = Ca, Sr or Ba) is nearly constant. tt E 10.27 State as to why (a) a solution of Na2CO3 is alkaline ? C (b) alkali metals are prepared by electrolysis of their fused chlorides ? (c) sodium is found to be more useful than potassium ? 10.28 Write balanced equations for reactions between no N (a) Na2O2 and water (b) KO2 and water (c) Na2O and CO2. 10.29 How would you explain the following observations? © (i) BeO is almost insoluble but BeSO4 is soluble in water, (ii) BaO is soluble but BaSO4 is insoluble in water, (iii) LiI is more soluble than KI in ethanol. 10.30 Which of the alkali metal is having least melting point ? (a) Na (b) K (c) Rb (d) Cs 10.31 Which one of the following alkali metals gives hydrated salts ? (a) Li (b) Na (c) K (d) Cs 10.32 Which one of the alkaline earth metal carbonates is thermally the most stable ? (a) MgCO3 (b) CaCO3 (c) SrCO3 (d) BaCO3 306 306 C:\ChemistryXI\Unit-10\Unit-10-Lay-3(reprint).pmd 27.7.6, C:\ChemistryXI\Unit-10\Unit-10-Lay-3(reprint).pmd Reprint16.10.6 27.7.6 (reprint)