s-Block Elements Notes

s-Block Elements Section (A) : General facts about elements The elements in which the last electron enters the outermost s-orbital are called s-block elements. As the s-orbital can accommodate only two electrons, two groups 1 and 2 belong to the s-block. Flame Test The alkali metals and alkaline earth metals and their salts impart characteristic colour to an oxidizing flame. Reason: This is because the heat from the flame excites the outermost orbital electron to a higher energy level. When they drop back to the ground state, there is emission of radiation in the visible region. The electrons in beryllium and magnesium are too strongly bound (due to small size) to get excited by flame. Hence, these elements do not impart any colour to the flame. Metal Colour Metal Colour Li Crimson red Be No colour Na Yellow Mg No colour K Violet / Lilac Ca Brick red Higher oxidation zone (melting zone) 1550ºC Higher reduction zone 1560ºC Rb Red violet Sr Crimson red Cs Blue Ba Apple green 520ºC Lower reduction zone 1450ºC Lower oxidation zone 300ºC Section (B) Based on Periodic trends Group – 1st(IA) Elements : (Alkali Metals) Atomic and Physical properties of the Alkali metals Lithium Sodium Potassium Rubidium Property Li Na K Rb Atomic number 3 11 19 37 Atomic mass (g mol–1) 6.94 22.99 39.10 85.47 Electronic configuration [He] 2s1 [Ne] 3s1 [Ar] 4s1 [Kr] 5s1 Ionization enthalpy / kJ mol–1 520 496 419 403 Hydration enthalpy/kJ mol–1 –506 –406 –330 –310 Metallic radius / pm 152 186 227 248 Ionic radius M+/ pm 76 102 138 152 m.p. / K 454 371 336 312 b.p / K 1615 1156 1032 961 –3 Density / g cm 0.53 0.97 0.86 1.53 –3.04 –2.714 –2.925 –2.930 Standard potentials E/ V for (M+/ M) Occurrence in lithosphere† 18* 2.27** 1.84** 78-12* *ppm (part per million), ** Percentage by weight Caesium Cs 55 132.91 [Xe] 6s1 376 –276 265 167 302 944 1.90 –2.927 2-6* Francium Fr 87 (223) [Rn] 7s1 ~375 – – (180) – – – – ~ 10–18* s-Block Elements Group IIA Elements (Alkaline Earth Metals) Property Atomic number Atomic mass (g mol–1) Electron configuration Ionization enthalpy (I) / kJ mol–1 Ionization enthalpy (II) /kJ mol–1 Hydration enthalpy (kJ/mol) Metallic radius / pm Ionic radius M2+/ pm m.p. / K b.p / K Density / g cm–3 Standard potential E / V for (M2+/ M) Occurrence in lithosphere Beryllium Be 4 9.01 [He] 2s2 899 1757 – 2494 112 31 1560 2745 1.84 –1.97 2* Magnesium Mg 12 24.31 [Ne] 3s2 737 1450 – 1921 160 72 924 1363 1.74 –2.36 2.76** Calcium Ca 20 40.08 [Ar] 4s2 590 1145 –1577 197 100 1124 1767 1.55 –2.84 4.6** Strontium Sr 38 87.62 [Kr] 5s2 549 1064 – 1443 215 118 1062 1655 2.63 –2.89 384* Barium Ba 56 137.33 [Xe] 6s2 503 965 – 1305 222 135 1002 2078 3.59 – 2.92 390* Radium Ra 88 226.03 [Rn] 7s2 509 979 – – 148 973 (1973) (5.5) –2.92 10–6* Section (C) & (D) : Based on Chemical Bonding, Properties of elements Properties of Alkali and Alkaline earth metals S.No. Atomic Properties Alkali metals 1. Outer Electronic configuration ns1 2. Oxidation number and valency (i)These elements easily form univalent +ve ion by losing loosely solitary ns 1 electron due to low IP value. 3. Atomic and Ionic radii Increase down the group, because value of n (principal quantum number) 4. Ionisation Energy As size increases, I.E. decreases down the group (so Cs have lowest I.P.) Order = Li > Na > K > Rb > Cs 5. Electropositive character or metallic character Alkali metals are strongly electropositive and metallic. Down the group electropositive nature increase so metallic nature also increases. i.e. M → M+ + e– Metallic Nature: Electropositive character 1 /I.P. Alkaline earth metals ns2 The IP1 of these metals are much lower than IP2 and thus it appears that these metals should form univalent ion rather than divalent ions but in actual practice, all these give bivalent ion. The atomic and ionic radii of the alkali earth metal are smaller than corresponding alkali metals. Reason higher nuclear charge (Zeff) On moving down the group size increase, as value of n increases. Down the group IE decreases due to increase in size. Be > Mg > Ca > Sr > Ba IE2 of Alkali metal > IE2 of Alkaline earth metal Reason IE1 of Alkaline earth metal is large due to increased nuclear charge in Alkaline earth metal as compared to Alkali metal but IE2 of Alkali metal is large because second electron in Alkali metal is to be removed from cation which has already acquired noble gas configuration Due to low IE they are strong electropositive but not as strong as Alkali metal because of comparatively high IE. The electropositive character increase down the group. 6. Hydration of ions 7. Photoelectric effect 8. Electronegativity S.No. Physical Property 1. Density 2. Hardness The cutting of sodium metal 3. 4. Melting points/ Boiling points Specific heat (i) Hydration represents for the dissolution of a substance in water to get absorb water molecule by weak valency forces Hydration of ions in the process when ions on dissolution in water get hydrated. (ii) Hydration energy  charge density on ion Degree of hydration  1/Cation size  charge  1/ionic mobility  1/conductivity Hydration energy = Li+ > Na+ > K+ > Rb+ > Cs+ (iii) Li+ being smallest in size has maximum degree of hydration and that is why lithium salts are mostly hydrated and moves very slowly under the influence of electric field. e.g : LiCl.2H2O. The phenomenon of emission of electrons when electromagnetic rays strikes against them is called photoelectric effect; Alkali metal have low I.P. so show photoelectric effect. Cs and K are used in Photoelectric cells. (i)These metals are highly electopositive and there by possess low values of electro negativities. (ii)Electronegativity of alkali metals decreases down the group. Order = Li > Na > K > Rb > Cs Hydration energy = Be2+ > Mg2+ > Ca2+ > Sr2+ >Ba2+ (i) Their electronegativities are also small but are higher than that of alkali metals (ii) Electronegativity decrease from Be to Ba Alkali metals Alkaline earth metals (i)All are light metals. (ii) Density increase down the group but K is lighter than Na. (i) All are silvery white metals. (ii) Light soft, malleable and ductile metals with metallic luster. (iii) Diamagnetic and colour less in form of ions. (iv) These metals are very soft and can be cut with a knife. Lithium is harder than any other alkali metal. The hardness depends upon cohensive energy. Cohensive energy  Force of attraction between atoms (i) Lattice energy decreases from Li to Cs and thus Melting points and Boiling points also decrease from Li to Cs. M.P. = Li > Na > K > Rb > Cs B.P. = Li > Na > K > Rb > Cs (i) Heavier than alkali metals. (ii) Density decrease slightly up to Ca after which it increases. (iii) Density of Mg is greater than Ca. Relatively soft but harder than Alkali metals. It decreases from Li to Cs. Li > Na > K > Rb > Cs (*need not to memorise) They have low Melting points and Boiling points but are higher than corresponding value of group I. Reason They have two valence electrons which may participate in metallic bonding compared with only one electron in Alkali metal. Consequently group II elements are harder and have higher cohesive energy and so, have much higher Melting points / Boiling points than Alkali metal . M.P. = Be > Ca > Sr >Ba > Mg ,B.P. = Be > Ba > Ca > Sr > Mg values are lesser than that of alkali metals, decreases down the group. *need not memorise. s-Block Elements S.No. Chemical Property 1. Alkali metals Alkaline earth metals (i) They generally form oxides and peroxides. O⎯ 2 M+O2 ⎯ M2O2 (Peroxide) ⎯→ M2O (Oxide) ⎯⎯→ The alkali metals tarnish in dry air due to the formation of their oxides on their surface. 4M + O2 —→ 2M2O They react vigorously in oxygen forming following oxides. 4 Li + O2 —→ 2 Li2O (Monoxide) 2 Na + O2 —→ Na2O2 (Peroxide) M + O2 —→ MO2 ( Superoxide) where M = K, Rb, Cs (i) Be and Mg are kinetically inert towards oxygen becasue of formation of a film of oxide on their surface. However powdered Be burn brilliantly. Principal Combustion Product (Minor Product) Metal Li Na Oxide Li2O (Na2O) Peroxide (Li2O2) Na2O2 KO2(Orange/Yellow Crystalline) RbO2 (Orange/Yellow Crystalline) CsO2 (Orange/Yellow Crystalline) K Rb Cs Alkali metals K Rb Cs  3Be + N2 (air) ⎯ Be3N2 ⎯→ (ii) Mg is more electropositive and burns with dazzling brilliance in air give MgO and Mg3N2.  Mg + O2(air) ⎯ MgO ; ⎯→  Mg + N2(air) ⎯ Mg3N2 ⎯→ Peroxides are coloured due to lattice defect. (Similar property with Li because both shows diagonal relation.) (iii) Ba gives BaO2 not BaO. (iv) Calcium, strontium and barium are readily attacked by air to form the oxide and nitride. They also react water with increasing vigour even in cold to form hydroxides. (v) BeO, MgO are used as refractory, because they have high M.P. (vi) Other metals (Ba or Sr form peroxide)  M + O2 ⎯ MO2 ⎯→ (i) Ca, Sr, Ba and Ra decompose cold water readily with evolution of hydrogen. M + 2H2O ⎯ ⎯→ M(OH) 2 + H2 (ii) Magnesium decomposes boiling water but beryllium is not attacked by water even at high temperatures as its oxidation potential is lower than the other members Fr – Standard potentials E/V for (M+/M) Na – 2.927 Li – 2.930 Property – 2.714 Action with water – 3.04 2. The oxides and peroxides are colourless when pure. (ii) All super oxide are paramagnetic and peroxides are diamagnetic in nature. (iii) The increasing stability of the peroxide or superoxide as the size of the metal ion increases is due to the stabilisation of large anions by larger cations through lattice energy effect. (iv) Since all the alkali metals are highly reactive towards air ; they are kept in kerosene oil. Reactivity increases from Li to Cs. (v) Only Lithium reacts with N2 (at room temperature) to form ionic lithium nitride Li3N because Li being strongest reducing agent converts N2 into N3–. 3Li + 1/2N2 —→ Li3N (i) Alkali metals decompose water to form the hydroxides having the formula MOH and dihydrogen. 2M + 2H2O —→ 2MOH(aq.) + H2(g) (M = An alkali metal). (ii) Li decompose water slowly, sodium reacts with water quickly K, Rb and Cs react with water vigorously. (iii) It may be noted that although lithium has most negative E value (In below table), its reaction with water is less vigorous than that of sodium which has the least negative E value among the alkali metals. This behaviour of lithium is attributed to its small size and very high hydration energy. It’s explanation lies in Kinetics, released energy in case of K, Rb, Cs is sufficient to melt or even vapourise and so more surface area is exposed to the water and kinetically reaction is faster than lithium. Other metals of the group react explosively with water. – 2.925 Action with O2 and N2 Superoxide  2Be + O2 (air) ⎯ 2BeO(amphoteric) ; ⎯→ (iv) They also react with proton donors such as alcohol, gaseous ammonia and terminal alkynes evolution of hydrogen. 2M + 2C2H5OH —→ 2C2H5OM + H2 Ethyl alcohol 3. Hydrides Metal ethoxide (i)They react with H2 forming metal hydride with formula MH which are of ionic nature. Stability of hydride decreases down the group. Since the electropositive character decreases from Cs to Li. 2M + H2 —→ 2MH (iii) The metal hydrides react with water to give MOH and H2. (act as reducing agent) (i) Except Be, all alkaline earth metals form hydrides (MH2) on heating directly with H2. (ii)The stability of hydrides decreases from Be to Ra. (iii) BeH2 is prepared by the action of LiAlH4 on BeCl2. BeCl2 +LiAlH4 —→ 2BeH2 + LiCl+ AlCl3 BeH2 & MgH2 is covalent and polymeric but other are s-Block Elements MH + H2O —→ MOH + H2 ionic. H Be Be H 4. (i)The alkali metals react vigorously with halogens to form ionic halides M+X–. 2M + X2 ⎯→ 2 M+X– (ii) Alkali metals halides (Cl2, Br2, I2) formation is increases form Li to Cs due to increase in electropositive character. Order of reactivity towards F2 Li > Na > K > Rb > Cs (iii) LiX have more covalent character (It is because of the high polarisation capability of Lithium ion (fajan’s rules)). (iv)Halides having ionic nature have high melting point and are good conductor of current in fused state. These are readily soluble in water. (v) Halides of potassium, rubidium and ceasium have property of combining with extra halogen atoms forming polyhalides. KI + I2 —→ KI3 H H H H (iv) The ionic hydrides of Ca, Sr, Ba liberate H2 at anode and metal at cathode. (i)The alkaline earth metals directly combine with halogens on heating to give metal halides MX2 (X=F,Cl,Br,I) (ii) Thermal decomposition of (NH4)2BeF4 is the best route for the preparation of BeF2, and BeCl2 is conveniently made from the oxide. BeO + C + Cl BeCl2 + CO Anhydrous beryllium halide can not be obtained from materials made in aqueous solution because the hydrated ions [Be(H2O)4]2+ is formed. i.e. [Be(H2O)4]Cl2 On dehydration, hydrolysis takes place. heat [Be(H2O)4]Cl2 ⎯⎯⎯ → Be(OH)2 + 2HCl (iii) Except for beryllium halides, all other halides of alkaline earth metals are ionic in nature. Beryllium halides are essentially covalent and soluble in organic solvents. Beryllium chloride has a chain structure in the solid state as shown below : Cl–Be–Cl Cl Halides Reducing nature (*need not to memorise) Fr – Cs – 2.927 Alkali metals K Rb – 2.930 Na – 2.925 Standard potentials E/V for (M+/M) Li – 2.714 Property – 3.04 5. (i) Reducing agent is electron donor. The alkali metals are strong reducing agents, lithium being the most and sodium the least powerful (above table). The standard electrode potential (E) which measures the reducing power represents the overall change : M(s) ⎯→ M(g) Sublimation enthalpy M(g) ⎯→ M+(g) + e– Ionization enthalpy M+(g) + H2O ⎯→ M+ (aq) Hydration enthalpy (ii) Lithium is expected to be least reducing agent due Cl Be Be Be Cl Cl In the vapour phase BeCl2 tends to form a chlorobridged dimer which dissociates into the linear monomer at high temperatures of the order of 1200 K. (iv)The ionic character of halides increases from Be to Ra. (v)Beryllium halides have covalent character due to small size and high effective nuclear charge and thus do not conduct electricity in molten state. (vi) The fluorides are relatively less soluble than the chlorides owing to their high lattice energies. (vii)The decreases in solubility of halides down the group is due to decrease in hydration energy because of increasing size of metal cation . (viii) The tendency to form halide hydrates gradually decreases (for example, MgCl2·6H2O, CaCl2·6H2O, SrCl2·6H2O and BaCl2·2H2O) down the group. The dehydration of hydrated chlorides, bromides and iodides of Ca, Sr and Ba can be achieved on heating; however, the corresponding hydrated halides of Be and Mg on heating suffer hydrolysis. (ix) CaCl2 has strong affinity with water and is used as dehydrating agent. The alkaline earth metals are strong reducing agents. This is indicated by large negative values of their reduction potentials (below table). However their reducing power is less than those of their corresponding alkali metals. Beryllium has less negative value compared to other alkaline earth metals. However, its reducing nature is due to large hydration energy associated with the small size of Be2+ ion and relatively large value of the atomization enthalpy of the metal. Property Alkaline earth metals Be Mg Ca Sr Ba Ra 6. Basic nature of hydroxide 7. Carbonates and bicarbonates (i) These oxides are easily hydrolysed by water to form the hydroxides. Thus M2O (oxide) + H2O ⎯→ MOH M2O2 (peroxide) + H2O ⎯→ 2 MOH +H2O2 MO2 (superoxide) + H2O ⎯→ 2 MOH +H2O2+O2 (ii) The Hydroxide which are obtained by the reaction of the oxide. With water all are white crystalline solids. The alkali metal hydroxides are the strongest of all bases and dissolve freely in water with evolution of much heat an account of intense hydration. Basic nature/Solubility in water/Thermal stability (i) The carbonates (M2CO3) and bicarbonates (MHCO3) are highly stable to heat, where M as alkali metals. (ii) Group 1 metals are so strongly basic, they (except lithium) also form solid bicarbonates. No other metals form solid bicarbonates. Lithium carbonate is not so stable to heat. Its hydrogencarbonate does not exist as a solid. Although NH4HCO3 also exists as a solid. (iii) The stability of these salts increases with the increasing electropositive character from Li to Cs. It is therefore Li2CO3 decompose on heating. Thermal stability/Solubility in water. LiHCO3 does not exist in solid form due to high polarizing power of Li+ and uncomparable size of Li+ cation and HCO3– anion.  Li2CO3 ⎯⎯→ Li2O + CO2 (iv) Bicarbonates are decomposed at relatively low temperature. – 2.92 – 2.92 – 2.89 – 2.84 Standard potentials E/V for (M+/M) – 2.36 to it's very high I.E. However, lithium has the highest hydration enthalpy which accounts for its high negative E value and its high reducing power. Reducing Nature in gas phase Reducing Nature in aqueous condition = Li > Cs > Rb > K > Na. – 1.97 s-Block Elements Basic/thermal stability (i) All these metal carbonates MCO3 are insoluble in neutral medium but soluble in acids and decompose on red heating. (ii) The stability of carbonates increases with increase in electropositive character of metal. (iii) Bicarbonates of alkaline earth metals do not exist in solid state but are known in solution only on heating their solution bicarbonates decomposed to liberate CO2 .  M(HCO3)2 ⎯⎯→ MCO3+CO2 + H2O (Solution) (iv)Solubility of carbonates decrease on moving down the group. BeCO3 > MgCO3 > CaCO3 > SrCO3 > BaCO3 2MHCO3 ⎯⎯⎯→ M2CO3 + H2O+ CO2 (v) Hydrolysis of carbonate Na2CO3 + 2H2O —→ 2NaOH + H2CO3 Li2CO3 + 2H2O —→ sparingly soluble (vi) The crystal structures of NaHCO3 and KHCO3 both show hydrogen bonding, but are different. (a) In NaHCO3, the HCO3– ions are linked into an infinite chain. (b) in KHCO3, RbHCO3, CsHCO3, HCO3– forms a dimeric anion. CsHCO3 300ºC (a) (b) 8. Complex ion formation 9. 10. Reaction with acids Formation of amalgams (i) A metal shows complex formation only when it has following characteristics. (a) Small size, (b) High nuclear charge, (c) Presence of empty orbitals in order to accept electron pair from ligand (electron pair donor species). (ii) Due to small size only Lithium in alkali metals, forms a few complex ions. Rest all alkali metals do not possess the tendency to form complex ion. Reacts vigorously with acids 2M + H2SO4 —→ M2SO4 + H2 (i) Alkali metals get dissolved in mercury to form amalgams with evolution of heat and the amalgamation is highly exothermic. Be2+ on account of smaller size forms many complexes such as [Be F3]–, [BeF4]2– Chlorophyll contains Mg2+ [Photosynthetic pigment in plants] (C.No.= 4) [Be(H2O)4]2+ + H2O —→ [Be(H2O)3OH]+ + H3O+ The alkaline earth metals readily react with acids liberated dihydrogen. M + 2HCl —→ MCl2 + H2 Alkaline earth metals get dissolved in mercury to form amalgams with evolution of heat and the amalgamation is highly exothermic. s-Block Elements (ii) Alkali metals form alloys themselves as well as with other metals. 11. Sulphates (i) All these form sulphates of type M2SO4. (ii) Except Li2SO4 rest all are soluble in water. Thermal stability /solubility in water (iii)These sulphates on fusing with carbon form sulphides. M2SO4 + 4C —→ M2S + 4CO 12. Sulphides All metals react with S forming sulphides such as Na2S and Na2Sn (n = 2, 3, 4, 5 or 6). The polysulphide ions are made from zig-zag chains of sulphur atoms. (i) Nitrates of both are soluble in water and decompose on heating. (ii) LiNO3 decomposes to give NO2 and O2 and rest all give nitrites and oxygen. 2MNO3 —→ 2MNO2 +O2 (except Li) 4LiNO3 —→ 2Li2O + 4NO2 + O2 13. Nitrates 2NaNO3 2NaNO3 (i) MSO4 type sulphates are formed (ii)The solubility of sulphates decreases on moving down the group. The sulphates of the alkaline earth metals are all white solids and stable to heat. BeSO4, and MgSO4 are readily soluble in water; the solubility decreases from CaSO4 to BaSO4. The greater hydration enthalpies of Be2+ and Mg2+ ions overcome the lattice enthalpy factor and therefore their sulphates are soluble in water. Thermal stability Solubility in water BeSO4 > MgSO4 > CaSO4 > SrSO4 > BaSO4 (iii) MSO4 + 2C —→ MS + 2CO2 M2+ + S2– —→ MS On heating they decompose into their corresponding oxides with evolution of a mixture of nitrogen dioxide and oxygen. ⎯→ MO + 2NO2 + ½O2 M(NO3)2 ⎯ (M = Be, Mg , Cr, Sr, Ba)  ⎯⎯→ 2NaNO2 + O2 500ºC  ⎯⎯→ Na2O + N2 + O2 800ºC Na 14. Nitride 15. Carbide 2NaNO3 ⎯⎯⎯ → Na2O + N2 + O2 Li3N + 3H2O —→ 3LiOH + NH3  When Li is heated with carbon, an ionic carbide Li2C2 is formed. 2Li + 2C —→ Li2C2 Other metals do not react with carbon directly but form carbides when heated with ethyne, or when ethyne is passed through a solution of metal in liquid ammonia. Na + C2H2 —→ NaH + C2 —→ Na2C2 [CC–H]– [CC]2– Na2C2 + 2H2O —→ NaOH + C2H2 Be3N2 + 6H2O ⎯→ 3Be(OH)2 + 2NH3  Mg3N2 + 6H2O ⎯→ 3Mg(OH)2 + 2NH3  The binary compounds of carbon with other elements (less electronegative or of similar electronegativity) are called carbides. They are classified into following 3 categories : (i) Ionic (ii) Covalent (iii) Interstitial (or metallic) (i) Ionic carbides (or salt like carbides) : Generally formed by the most electropositive elements such as alkali and alkaline earth metals and aluminium (Boron is exception). Based on the product obtained on hydrolysis, they are further sub-classified into three types. (a) Methanides : These give CH4 on reaction with H2O. Al4C3 + 12H2O —→ 4Al(OH)3 + 3CH4 ; Be2C + 4H2O —→ 2Be(OH)2 + CH4 These carbides contain C4– ions in their constitution. (b) Acetylides : These give C2H2 on reaction with H2O. CaC2 + 2H2O —→ Ca(OH) 2 + C2H2 Al2 (C2)3 + 6H2O —→ 2Al(OH)3 + 3C2H2 SrC2 + 2H2O —→ Sr (OH) 2 + C2H2 Such compounds contain C22– ions. (c) Allylides : These give 1-propyne on reaction with H2O. Mg2C3 + 4H2O —→ 2Mg(OH)2 + CH3–CCH .. Such compounds contain C34– [: C – C  C :]4 − ions. .. (ii) Covalent carbides : Molecules like SiC and B4C are also examples of covalent carbides. (iii) Interstitial or metallic carbides Such carbides are formed by transition metals and some of the lanthanides and actinides. Interstitial carbides retain many of the properties of metals. They conduct electricity by metallic conduction and have properties of metals (a lusture like a metal). In these compounds carbon atoms occupy octahedral holes in the closed packed metal lattice. These are s-Block Elements generally very hard and have very high melting point (e.g. WC). Carbides of Cr, Mn, Fe, Co and Ni are hydrolysed by water or dilute acids. Lattice Energy: Energy change when one mole of crystalline lattice is formed from gaseous ions eg. 2Al3+ + 3O2– —→ Al2O3 + L.E. Hydration Energy: It is the energy change when gaseous ions form aqueous ions. eg. Na+ + aq. —→ Na+ + H.E. of Na+ SO42- + aq. —→ SO42- + H.E. of SO42– Solutions in liquid NH3 Alkali metals dissolve in liquid ammonia (high conc. 3 M) and give blue solution which is conducting, reducing and paramagnetic in nature. Reason On dissolving Metal in NH3 M(s) + NH3( ) ⎯→ M+(NH3) + e–(NH3) M+ + x (NH3) ⎯→ [M(NH3)x]+ ⎯→ Ammoniated cation e– + y (NH3) ⎯→ [e(NH3)y]– ⎯→ Ammoniated electron The blue colour is due to ⎯→ Ammoniated electron The paramagnetic nature is due to ⎯→ Ammoniated electron The conducting nature is due to ⎯→ Ammoniated M+ + Ammoniated electron On standing the colour fades due to formation of amide after liberating hydrogen. M+ + e– + NH3 ⎯→ MNH2(amide) + H2(g) In the absence of impurities like. Fe, Pt, Zn etc, the solutions are stable. In concentrated solution, the blue colour changes to bronze colour and diamagnetic due to the formation of metal clusters and ammoniated electrons also associate to form electron pairs 2e– (NH3)y ⎯→ [ e–(NH3)y]2 Solutions are of much lower density than the pure solvent, i.e., they occupy for greater volume than that expected from the sum of the volumes of metal and solvent Peroxide and superoxides of Na & K are widely used as oxidising agent and air purifiers in space capsules, submarines and breathing mask. Alkaline metal in liq. NH3 Like alkali metals, the alkaline earth metals dissolve in liquid ammonia to give deep blue black solutions forming ammoniated ions M + (x + y)NH3 —→ [M(NH3)x]2+ + 2[e(NH3)y]– From these solutions, the hexa-ammoniates [M(NH3)6]2+ can be recovered. Uses of alkali metal (1) Lithium metal is used to make useful alloys, • with lead to make ‘white metal’ bearings for motor engines. • with aluminium to make aircraft parts. • with magnesium to make armour plates. (2) It is used in thermonuclear reactions. (3) Lithium is also used to make electrochemical cells. (4) Sodium is used to make a Na/Pb alloy needed to make PbEt4 and PbMe4. These organolead compounds were earlier used as anti-knock additives to petrol, but nowadays vehicles use lead-free petrol. (5) Liquid sodium metal is used as a coolant in fast breeder nuclear reactors. (6) Potassium chloride is used as a fertilizer. (7) Potassium hydroxide is used in the manufacture of soft soap. It is also used as an excellent absorbent of carbon dioxide. (8) Caesium is used in devising photoelectric cells. Uses of alkaline metal (1) Beryllium is used in the manufacture of alloys. (2) Copper-beryllium alloys are used in the preparation of high strength springs. (3) Metallic beryllium is used for making windows of X-ray tubes. s-Block Elements (4) Magnesium forms alloys with aluminium, zinc, manganese and tin. Magnesium-aluminium alloys being light in mass are used in air-craft construction. (5) Magnesium (powder and ribbon) is used in flash powders and bulbs, incendiary bombs and signals. (6) A suspension of magnesium hydroxide in water (called milk of magnesia) is used as antacid in medicine. (7) Magnesium carbonate is an ingredient of toothpaste. (8) Calcium is used in the extraction of metals from oxides which are difficult to reduce with carbon. (9) Calcium and barium metals, owing to their reactivity with oxygen and nitrogen at elevated temperatures, have often been used to remove air from vacuum tubes. (10) Radium salts are used in radiotherapy, for example, in the treatment of cancer. Biological Importance Of Sodium And Potassium: • Sodium ions are found primarily on the outside of cells, being located in blood plasma and in the interstitial fluid which surrounds the cells. These ions participiate in the transmission of nerve signals, in regulating the flow of water across cell membranes and in the tranasport of sugars and amino acids into cells. Sodium and potassium, although so similar chemically, differ quantitatively in their ability to penetrate cell membranes, in their transport mechanisms and in their efficiency to activate enzymes. Thus, potassium ions are the most aundant cations within cell fluids, where they activate many enzymes, participate in the oxidation of glucose to produce ATP and, with sodium, are responsible for the transmission of nerve signals A typical 70 kg man contains about 90 g of Na and 170 g of K compared with only 5 g of iron and 0.06 g of copper. Biological Importance of Magnesium and Calcium : • Monovalent sodium and potassium ions and divalent magnesium and calcium ions are found in large proportions in biological fluids. These ions perform important biological functions such as maintenance of ion balance and nerve impulse conduction. • All enzymes that utilise ATP in phosphate transfer require magnesium as the cofactor. The main pigment for the absorption of light in plants is chlorophyll which contains magnesium. About 99 % of body calcium is present in bones and teeth. It also plays important roles in neuromuscular function, interneuronal transmission, cell membrane integrity and blood coagulation. • The calcium concentration in plasma is regulated at about 100 mgL–1. It is maintained by two hormones : calcitonin and parathyroid hormone. Do you know that bone is not an inert and unchanging substance but is continuously being solubilised and redeposited to the extent of 400 mg per day in man? All this calcium passes through the plasma. An adult body contains about 25 g of Mg and 1200 g of Ca compared with only 5 g of iron and 0.06 g of copper. The daily requirement in the human body has been estimated to be 200–300 mg. ANOMALOUS PROPERTIES OF LITHIUM The anomalous behavior of lithium is due to the : (i) Exceptionally small size of its atom and ion, (ii) High polarising power (i.e., charge/ radius ratio). As a result, there is increased covalent character of lithium compound which is responsible for their solubility in organic solvent. Further, lithium shows diagonal relationship to magnesium. S.No. 1. Property Hardness Li is much harder. 4. 5. M.P and B.P Reactivity Reducing agent Combustion in air Higher M.P and B.P Less reactive Strong Li form monoxide (Li2O) and nitride (Li3N) ; not for other. 6. Hydration of ion 2. 3. 8. Hydrogen Carbonate Ethynide 9. Lithium nitrate 7. Li Favored for Li ; not for other. Li has maximum degree of hydration for this reason. Lithium salts are mostly hydrated. E.g LiCl.2H2O. Li is not obtained in the solid form while all other elements form solid hydrogen carbonates. Favored for Li ; not for other. 4LiNO  3 ⎯⎯→ 2Li2O + 4NO2 + O2 Lithium Oxide s-Block Elements 10. LiF and Li2O 11. Carbide 12. Hydroxide 13. 14. Carbonate 15. Bicarbonate Nitrite Complex ion formation Reaction with NH3 16. 17. Where as other alkali metal nitrates decompose to give the corresponding nitrite.  2NaNO3 ⎯⎯→ 2NaNO2 + O2 Sodium nitrite These are much less soluble in water. Solubility in water is less than the corresponding compounds of other alkali metal. Li reacts directly with carbon to form anionic carbide. Lithium hydroxide is less basic Li2CO3, LiNO3 and LiOH all form the oxides on gentle heating. Less stable. Less stable. Lithium forms a bicarbonates in solution it does not form a solid bicarbonate. Where as the other all forms stable solid bicarbonates. Lithium has a great tendency to form. Complexes not for other. Due to small size of Lithium. Li when heated in NH3 imide (Li2NH) while other alkali metals form amides (MNH2) Points of Similarities between Lithium and Magnesium The similarity between lithium and magnesium is particularly striking and arises because of their similar size: atomic radii, Li = 152 pm, Mg = 160 pm; ionic radii : Li + = 76 pm, Mg2+ = 72 pm. The main points of similarity are : S.No. Properties Li and Mg 1. Hardness Li and Mg are much harder. 2. Density These are lighter than other elements in the respective group. 3. Reaction with water Both react slowly with water. 4. Solubility of hydroxide and oxide Less soluble and their hydroxides decompose in acid on heating. 5. Reaction with N2 By direct combination with nitrogen both form a nitride Li 3N and Mg3N2. 6. Oxides The oxides Li2O and MgO donot combine with excess oxygen to give any superoxide. 7. Carbonates Carbonates of both decompose easily on heating to form the oxides and CO2. Solid hydrogen carbonates are not formed by Li and Mg. 8. Solubility of halides in ethanol Both LiCl and MgCl2 are soluble. 9. Hydration of ion Both LiCl and MgCl2 are deliquescent and crystallise from aqueous solution as hydrates, LiCl.H2O and MgCl2.6H2O. Anamolous Behaviour of Beryllium The properties of beryllium the first member of the alkaline earth metal, differ from the rest of the member. Its is mainly because of (i) Its small size and high polarizing power. (ii) Relatively high electro negativity and ionization energy as compared to other members. (iii) Absence of vacant d–orbitals in its valence shell. Some important points of difference between beryllium and other members (especially magnesium) are given below. s-Block Elements S.No. Properties 1. Hardness Be is harder than other members of its group 2. 3. 4. 5. 6. Density M.P. and B.P. Reaction with water Nature of oxides Nature of compounds Be is lighter than Mg Higher then other members of its group. Be does not react with water while Mg reacts with boiling water. BeO is amphoteric while MgO is weakly basic. Be forms covalent compounds whereas other members form ionic compounds. 7. Carbide Beryllium carbide reacts with water to give methane whereas carbides of other alkaline earth metals gives acetylene gas. Be2C + 4H2O → 2Be (OH)2 + CH4 MgC2 + 2H2O → Mg (OH)2 + C2H2 CaC2 + 2H2O → Ca (OH)2 + C2H2 8. Hydride The beryllium hydride is electron deficient and polymeric, with muti center bonding like aluminium hydride. 9. Co-ordination number Beryllium does not exhibit coordination number more than four as it has four orbitals in the valence shell. The other members of this group has coordination number 6. 10. Reaction with Alkali Be dissolves in alkalies with evolution of hydrogen Be + 2NaOH +2H2O→ Na2BeO2.2H2O + H2 (sodium beryllate Other alkaline earth metals don't react with alkalies. Resemblance of Beryllium with Aluminium (Diagonal relationship) The following points illustrate the anomalous behaviour of Be and its resemblance with Al. S.No. Properties Be and Al 1. 2. Nature of compounds Nature of hydroxide Unlike groups-2 elements but like aluminium, beryllium forms covalent compounds. The hydroxides of Be, [Be(OH)2] and aluminium [Al(OH)3] are amphoteric in nature, whereas those of other elements of group – 2 are basic in nature. 3. 4. Nature of oxide Polymeric structure The oxides of both Be and Al i.e. BeO and Al2O3 are high melting insoluble solids. BeCl2 and AlCl3 have bridged chloride polymeric structure. 5. 6. Salts Carbides The salts of beryllium as well as aluminium are extensively hydrolysed. Carbides of both the metal reacts with water liberating methane gas. Be2C + 4H2O → 2Be (OH)2 + CH4 Al4C3 + 12H2O → 4Al(OH)3 + 3CH4 7. Oxides and hydroxides The oxides and hydroxides of both Be and Al are amphoteric and dissolve in sodium hydroxide as well as in hydrochloric acid. BeO + 2HCl → BeCl2 + H2O BeO + 2NaOH → Na2BeO2 + H2O Al2O3 + 6HCl → 2AlCl3 + H2O Al2O3 + 2NaOH → 2NaAlO2 + H2O 8. Reaction with acids Like Al, Be is not readily attacked by acids because of the presence of an oxide film. s-Block Elements Section (E) : Oxides, Peroxides, Super Oxides, Hydroxides GROUP -I & II OXIDES 1. Sodium Oxide (Na2O) Preparation: Name of compound Name and Brief about the process Related chemical reaction (1) By burning sodium at 180°C in a limited supply of air or oxygen and distilling off the excess of sodium in vacuum. Sodium Oxide (Na2O) (2) By heating sodium peroxide, nitrate or nitrite with sodium. (3) Sodium oxide is formed when the mixture of sodium azide and sodium nitrite is heated. 1 180º → Na2O O2 ⎯⎯⎯ 2 Na2O2 + 2Na —→ 2Na2O 2NaNO3 + 10Na —→ 6Na2O + N2 2NaNO2 + 6Na —→ 4Na2O + N2 2Na + 3NaN3 + NaNO2 ⎯→ 2Na2O + 5N2 Chemical Properties: It is white amorphous substance. It dissolve violently in water, yielding caustic soda (NaOH) and evolving a large amount of heat. Na2O + H2O ⎯→ 2NaOH Uses : It is used as dehydrating and polymerising agent in organic chemistry. (1) (2) 2. (1) (2) (3) (4) (5) (6) (1) (2) (3) Sodium Peroxide (Na2O2) Preparation Name of compound Name and Brief about the process Related chemical reaction 2Na + O2 (excess) ⎯⎯⎯→ Na2O2 Sodium Peroxides (Na2O2) (1) By heating the metal in excess of air or oxygen at 300°, which is free from moisture and CO2. (2) Industrial method : It is a two stage reaction in the presence of excess air. 300ºC 2Na + O2 —→ Na2O Na2O + O2 —→ Na2O2 Properties: It is a pale yellow solid (when impure), becoming white in air from the formation of a film of NaOH and Na2CO3. In cold water (~0°C) produces H2O2 but at room temperature produces O2. In ice-cold mineral acids also produces H2O2. ~0ºC → 2NaOH + H2O2 Na2O2 + 2H2O ⎯⎯⎯ 25ºC 2Na2O2 + 2H2O ⎯⎯⎯→ 4NaOH + O2 ~0ºC → Na2SO4 + H2O2 Na2O2 + H2SO4 ⎯⎯⎯ 25ºC 2Na2O2 + H2SO4 ⎯⎯⎯→ 2Na2SO4 + 2H2O + O2 It reacts with CO2, giving sodium carbonate and oxygen and hence its use for purifying air in a confined space e.g. submarine, ill-ventilated room. 2Na2O2 + 2CO2 —→ 2Na2CO3 + O2 Na2O2 + CO —→ Na2CO3 It is an oxidising agent and oxidises charcoal, CO, NH 3, SO2. 3Na2O2 + 2C —→ 2Na2CO3 + 2Na [deposition of metallic Na] CO + Na2O2 —→ Na2CO3 SO2 + Na2O2 —→ Na2SO4 2NH3 + 3Na2O2 —→ 6NaOH + N2 Sulphides are oxidised to corresponding sulphates Na2O2 —→ Na2O + [O] ; Na2S + 4[O] —→ Na2SO4 Na2O2 —→ Na2O + [O]; 2Al + 3[O] —→ Al2O3; Al2O3 + Na2O —→ 2NaAlO2. Uses : For preparing H2O2, O2. Oxygenating the air in submarines. Oxidising agent in the laboratory. Oxides of Potassium K2O Colours White K2O2 K2O3* KO2 KO3 White Red Bright Yellow Orange Red Solid s-Block Elements 3. Potassium sesquioxide (need not memorize). Preparation: Name of compound Name and Brief about the process Related chemical reaction ⎯heating ⎯⎯→ heating ⎯⎯⎯→ K2O 2KNO3 + 10K K2O (Potassium oxide) By heating potassium. potassium nitrate with K2O (White) 6K2O + N2 (Yellow) K2O + H2O —→ 2KOH Name of compound K2O2 (Potassium peroxide) Name of compound KO2 (Potassium superoxide) Name of compound K2O3 (Potassium sesquioxide) Name and Brief about the process By burning potassium at 300°C in a limited supply of air or oxygen. Name and Brief about the process Pas (i) Passage of O2 through a blue solution of K in liquid NH3 yields oxides K2O2 (white), K2O3 (red) and KO2 (deep yellow) i.e KO2 reacts with H2O and produces H2O2 and O2 both. (ii) It is prepared by burning potassium in excess of oxygen free from moisture. Related chemical reaction 2K + O2 ⎯⎯ ⎯⎯→ K2O2 Controlled air at 300 º C Related chemical reaction O2 K in liq. NH3 ⎯⎯→ K2O2 —→ K2O3 —→ KO2 white red yellow 0C 2KO2 + 2H2O ⎯~⎯ ⎯→ 2KOH + H2O2 + O2 K + O2 —→ KO2 Name and Brief about the process It is obtained when oxygen is passed through liquid ammonia containing potassium. Related chemical reaction 3O2 → 2K2O3 4K (dissolved in liquid NH3) ⎯⎯⎯ Name of compound Name and Brief About the process Related chemical reaction KO3 (Potassium ozonide) From KOH −10 to−15C KOH + O3 (ozonised oxygen) ⎯⎯⎯⎯⎯→ KO3 (Dry powdered) (orange solid) Properties of Potassium superoxide (KO2) It is a orange coloured (chrome yellow) powder and reacts with water according to following reaction. 2KO2 + 2H2O —→ 2KOH + H2O2 + O2 It reacts directly with CO and CO2. 2KO2 + CO —→ K2CO3 + O2 ; 2KO2 + CO2 —→ K2CO3 + O2 If more CO2, in presence of moisture is present; then 4KO2 + 4CO2 + 2H2O —→ 4KHCO3 + 3O2 On heating with sulphur, it forms potassium sulphate 2KO2 + S —→ K2SO4 Uses : It is used as an oxidising agent and air purifier in space capsules, submarine and breathing mask as it produces O2 and removes CO2. 4. Magnesium Oxide (MgO): Name of compound Name and Brief about the process It is also called magnesia and obtained by Magnesium Oxide (MgO) heating natural magnesite. (1) (2) (3) Related chemical reaction MgCO3 —→ MgO + CO2 Properties : It is white powder. It's m.p. is 2850°C. Hence used in manufacture of refractory bricks for furances. And it is acts as basic flux and facilitates the removal of acidic impurities of Si, P and S from steel through slag formation. It is very slightly soluble in water imparting alkaline reaction. s-Block Elements 5. Calcium Oxide (CaO): Preparation Name of compound * (1) Name and Brief about the process Related chemical reaction It is commonly called as quick lime or lime Calcium Oxide CaO + CO2 and made by decomposing lime stone at a CaCO3 (CaO) high temperature about 1000°C. The Carbon dioxide is removed as soon as it is produced to enable the reaction to proceed to completion. Chemical Properties : It is white amorphous powder of m.p. 2570°C. On exposure to atmosphere; it absorbs moisture and carbondioxide. CaO + H2O —→ Ca(OH) 2 ; CaO + CO2 —→ CaCO3 acidic oxide (2) (3) (4) It emits intense light (lime light), when heated in oxygen-hydrogen flame. It combines with limited amount of water to produce slaked lime. This process is called slaking of lime. Quick lime slaked with soda gives solid sodalime (CaO). Being a basic oxide. CaO + H2O —→ Ca(OH) 2 Soda lime (basic oxide) combines with some acidic oxides at high temperature. CaO + SiO2 —→ CaSiO3 6CaO + P4O10 —→ 2Ca3(PO4)2 Uses : It is an important primary material for manufacturing cement and is the cheapest form of alkali. It is used in the manufacture of sodium carbonate from caustic soda. It is employed in the purification of sugar and in the manufacture of dye stuffs. Magnesium Peroxide (MgO2) and Calcium Peroxide (CaO2) These are obtained by passing H2O2 in a suspension of Mg(OH) 2 and Ca(OH) 2. Uses : MgO2 is used as an antiseptic in tooth paste and as a bleaching agent. (i) (ii) (iii) HYDROXIDES 1. Sodium Hydroxides(Caustic Soda) NaOH (White) : Preparation : Name of compound Name and Brief about the process Related chemical reaction Cathode:Na++ e– ⎯⎯→ Na-amalgam Hg Anode : Cl– ⎯→ 1 Cl2 + e– 2 2Na-amalgam + 2H2O ⎯→ 2NaOH + 2Hg + H2 Sodium Hydroxides (NaOH) (1) Electrolysis of Brine : Sodium hydroxide is prepared by the electrolysis of sodium chloride in Castner-Kellner cell. A brine solution is electrolysed using a mercury cathode and a carbon anode. Sodium metal discharged at the cathode combines with mercury to form sodium amalgam. Chlorine gas is evolved at the anode. The amalgam is treated with water to give sodium hydroxide and hydrogen gas. Anode Cathode (2) By Diaphragm cell (3) Caustication of Na2CO3 (Gossage's method) Na2CO3 + Ca(OH)2 ⎯→ 2NaOH + CaCO3 (suspension) (1) (2) (3) (4) (5) (6) (7) (8) (9) Properties: Sodium hydroxide is a white, translucent solid. It melts at 591 K. It is readily soluble in water to give a strong alkaline solution. Crystals of sodium hydroxide are deliquescent. The sodium hydroxide solution at the surface reacts with the CO2 in the atmosphere to form Na2CO3. It is white crystalline, deliquescent, highly corrosive solid. It is stable towards heat. It's aqueous solution alkaline in nature and soapy in touch. NH4Cl + NaOH ⎯→ NaCl + NH3  + H2O FeCl3 + 3NaOH ⎯→ Fe(OH)3  + 3NaCl Brown ppt ZnCl2 + 2NaOH ⎯→ Zn(OH)2 + 2NaCl Zn(OH)2  + 2NaOH ⎯→ Na2ZnO2 + 2H2O [Same with AlCl3, SnCl2, PbCl2] soluble Acidic and amphoteric oxides gets dissolved easily e.g. CO2 + 2NaOH ⎯→ Na2CO3 + H2O Al2O3 + 2NaOH ⎯→ 2NaAlO2 + H2O Aluminium and Zn metal gives H2 from NaOH. 2Al + 2NaOH + 2H2O ⎯→ 3H2 + 2NaAlO2 Several non metals such as P, S, Cl etc. yield a hydride instead of hydrogen.e.g. 4P + 3NaOH + 3H2O ⎯→ PH3 + 3NaH2PO2 (Disproportionation reaction) NaOH is stable towards heat but reduced to metal when heated with carbon. 2NaOH + C ⎯→ 2Na + 2 CO + H2 +2 (10). NaOH + Metal Oxide (M) +3 +4 (i) (ii) (iii) (iv) (v) (vi) 2. (1) (2) (3) (4) (1) (2) (3) (4) Na2MO2 Na3MO3 Na2MO3 Above are general reactions of NaOH with metal oxides having metal's Oxidation number +2, +3 & +4 respectively. Uses : It is used in The manufacture of soap, paper, artificial silk and a number of chemicals. In petroleum refining. In the purification of bauxite. In the textile industries for mercerising cotton fabrics. For the preparation of pure fats and oils . As a laboratory reagent. Potassium Hydroxide (KOH): Preparation: It is prepared by electrolysis of KCl solution. KOH resembles NaOH in all its reactions. However KOH is much more soluble in alcohol. This accounts for the use of alcoholic KOH in organic chemistry. KOH is called caustic potash, because of their corrosive properties (for example on glass or on skin) and its aqueous solution is known as potash lye. 2KOH + 4NO ⎯→ 2KNO2 + N2O + H2O 4KOH + 6NO ⎯→ 4KNO2 + N2 + 2H2O It is used for the absorption of gases like CO2, SO2, etc. It is used for making soft soaps. Properties: Same as NaOH It is stronger base compared to NaOH. Solubility in water is more compared to NaOH. In alcohol, NaOH is sparingly soluble but KOH is highly soluble. As a reagent KOH is less frequently used but in absorption of CO 2, KOH is preferably used compared to NaOH. Because KHCO3 formed is soluble whereas NaHCO3 is insoluble and may therefore choke the tubes of apparatus used. Na2PbO3 Na2PbO2 +Na2SO4 s-Block Elements s-Block Elements 3. Magnesium Hydroxide (Mg(OH)2): It occurs in nature as the mineral brucite. Preparation: (1) (2) (3) Name of compound Name and Brief about the process Related chemical reaction Magnesium Hydroxide (Mg(OH)2) It can be prepared by adding caustic soda solution to a solution of Magnesium sulphate or chloride solution. MgSO4 + 2NaOH ⎯→ Mg(OH)2 + Na2SO4 MgCl2 + 2NaOH ⎯→ Mg(OH)2 + 2NaCl MgCl2 + Ca(OH)2 ⎯→ Mg(OH)2 + CaCl2 MgO + H2O ⎯→ Mg(OH)2 Chemical Properties: It can be dried at temperature upto 100°C only otherwise it breaks into its oxide at higher temperature. Mg(OH)2 ⎯→ MgO + H2O It is slightly soluble in water imparting alkalinity. It dissolves in NH4Cl solution. Mg(OH)2 + 2NH4Cl ⎯→ MgCl2 + 2NH3.H2O Thus, Mg(OH)2 is not therefore precipitated from a solution of Mg2+ ions by NH3.H2O. in presence of excess of NH4Cl. Uses : A suspension of Mg(OH)2 in water is used in medicine as an antacid (An antacid is substance which neutralizes stomach acidity) under the name, milk of magnesia. 4. Calcium Hydroxide (Ca(OH)2) (White Powder): Preparation : (1) (2) (3) (4) (5) (i) (ii) (iii) Name of compound Name and Brief about the process Related chemical reaction Calcium Hydroxide (Ca(OH)2) By spraying water on quicklime. CaO + H2O ⎯→ Ca(OH)2 Properties: It is a white amorphous powder. It is sparingly soluble in water. It's solubility in hot water is less than that of cold water. Hence solubility decreases with increase in temperature. The aqueous solution is known as lime water and a suspension of slaked lime in water is known as milk of lime. When carbon dioxide is passed through lime water it turns milky due to the formation of calcium carbonate. Ca(OH)2 + CO2 ⎯→ CaCO3 + H2O On passing excess of carbon dioxide, the precipitate dissolves to form calcium hydrogen carbonate. CaCO3 + CO2 + H2O ⎯→ Ca(HCO3)2 Milk of lime reacts with chlorine to form hypochlorite, a constituent of bleaching powder. 2Ca(OH) 2 + 2Cl2 ⎯→ CaCl2 + Ca(OCl)2 + H2O Bleaching powder Uses: It is used in the preparation of mortar, a building material. It is used in white wash due to its disinfectant nature. It is used in glass making, in tanning industry, for the preparation of bleaching powder and for purification of sugar. s-Block Elements Section (F) : Carbonates, Bicarbonates CARBONATES 1. Sodium Carbonate (Washing soda) Na2CO3.10H2O (White Solid) : Preparation: 1. Name of compound Name and Brief about the process Related chemical reaction mild heating NaCl + H2SO4(conc.) ⎯⎯ ⎯ ⎯⎯→ NaHSO4 + HCl NaCl + NaHSO4 ⎯⎯ ⎯ ⎯→ Na2SO4 + HCl Strongly heated (1) Leblanc Process Sodium Carbonate (Washing soda) Na2CO3.10H2O (Salt Cake) Na2SO4 + 4C ⎯→ Na2S + 4CO  Na2S + CaCO3 ⎯→ Na2CO3 + CaS (2) Solvay Process Step-1 (In ammonia absorber) (i)Saturation of brine with ammonia and CO2 (ii) Ammoniated brine is filtered to remove calcium and magnesium impurities as their insoluble carbonates. Step-2 (In carbonation tower) : (i) Formation of insoluble NaHCO3 (ii) Reaction is exothermic and hence there is a cooling arrangement. (iii) NaHCO3 is insoluble in cold brine solution because of the common ion effect. It is separated by filtration and the filtered is used for recovering NH3 & CO2. Step-3 (Calcination to get sodium carbonate) : Step - 4 (In recovery tower) : Recovery of ammonia and carbondioxide. CaCl2 is obtained as by product. 2NH3 + CO2 + H2O ⎯→ (NH4)2CO3 CaCl2 + (NH4)2CO3 ⎯→ CaCO3  + 2NH4Cl MgCl2 + (NH4)2CO3 ⎯→ MgCO3  + 2NH4Cl NH3 + CO2 + H2O ⎯→ NH4HCO3 ; 30 º C NH4HCO3 + NaCl ⎯⎯ ⎯→ NaHCO3 + NH4Cl 150 º C 2 NaHCO3 ⎯⎯ ⎯→ Na2CO3 + CO2 + H2O  / Steam NH4 HCO3 ⎯⎯ ⎯⎯→ NH2 + CO2 + H2O  / Steam 2NH4 Cl +Ca(OH)2 ⎯⎯ ⎯⎯→ 2NH3+ 2H2O+CaCl2 * advantage is taken of low solubility of NaHCO3, it gets precipitated in the reaction of NaCl + NH4HCO3. 2. Naturally from trona heat 2(Na2CO3.NaHCO3.2H2O) ⎯⎯⎯ → 3Na2CO3 + CO2 + 5H2O (1) (2) (3) (4) Properties Anhydrous Na2CO3 is called as soda ash, which does not decompose on heating but melts at 852°C. Sodium carbonate is a white crystalline solid which exists as a decahydrate, Na 2CO3·10H2O. This is also called washing soda. It is readily soluble in water. On heating, the decahydrate loses its water of crystallisation to form monohydrate. Above 373K, the monohydrate becomes completely anhydrous and changes to a white powder called soda ash. 375 K ⎯→ Na2CO3·H2O + 9H2O Na2CO3·10H2O ⎯⎯  373 K ⎯ ⎯→ Na2CO3 + H2O Na2CO3·H2O ⎯⎯ (soda ash) Carbonate part of sodium carbonate gets hydrolysed by water to form an alkaline solution. Na2CO3 + H2O ⎯→ H2CO3 (weak acid) + NaOH (strong) Na2CO3 absorbs CO2 yielding sparingly soluble sodium bicarbonate which can be calcined at 250° to get pure sodium carbonate. Na2CO3 + H2O + CO2 2NaHCO3 (solid) It dissolved in acid with effervescence of CO2 and causticised by lime to give caustic soda. Na2CO3 + HCl ⎯→ 2NaCl + H2O + CO2 Na2CO3 + Ca(OH)2 ⎯→ 2NaOH + CaCO3 s-Block Elements Uses : (i) (ii) (iii) (iv) It is used in water softening, laundering and cleaning. It is used in the manufacture of glass, soap, borax and caustic soda. It is used in paper, paints and textile industries. It is an important laboratory reagent both in qualitative and quantitative analysis. 2. Potassium Carbonate (K2CO3) Name of compound Potassium Carbonate (K2CO3) Name and Brief about the process By leblance process, it can be prepared but by solvay process it cannot be prepared because KHCO3 is soluble in water. Related chemical reaction mild heating KCl + H2SO4 (conc.) ⎯⎯ ⎯ ⎯⎯→ KHSO4 + HCl KCl + KHSO4 ⎯⎯ ⎯ ⎯→ K2SO4 + HCl Strongly heated K2SO4 + 4C ⎯→ K2S + 4CO  K2S + CaCO3 ⎯→ K2CO3 + CaS Properties: It resembles with Na2CO3, m.p. is 900°C but a mixture of Na2CO3 and K2CO3 melts at 712°C. Uses It is used in glass manufacturing. * need not memories. Note : Calcium carbonate and Magnesium carbonate found in nature. Calcium bicarbonate and Magnesium bicarbonate are present in temporary hardness of water. Unstable and unimportant. Same for KHCO3. Section (G) : Chlorides, Sulphates CHLORIDES Sodium Chloride (NaCl) and Potassium Chloride, Calcium Chloride Preparation: NaCl : Found in nature as rock salt or in sea water. KCl : Found in nature as sylvine (KCl) or carnallite (2KCl.MgCl2.6H2O) CaCl2 : Obtained as byproduct in Solvay’s process. (3) Properties of NaCl : It is nonhygroscopic but the presence of MgCl2 in common salt renders it hygroscopic. It is used to prepare freezing mixture in laboratory [Ice-common salt mixture is called freezing mixture and temperature goes down to –23°C.] For melting ice and snow on road. (i) (ii) Uses of NaCl: It is used as a common salt or table salt for domestic purpose. It is used for the preparation of Na2O2, NaOH and Na2CO3. (1) (2) Magnesium Chloride (MgCl2) It occurs in nature as mineral carnallite, KCl.MgCl2.6H2O. (1) (2) (3) (4) Preparation : By Dow’s Processes (Natural Brine process and Dolomite process). See Metallurgy, stdXII. Properties: It crystallises as hexahydrate. MgCl2. 6H2O It is deliquescent solid. This hydrate undergoes hydrolysis as follows: MgCl2·6H2O ⎯→ Mg(OH)Cl + HCl + 5H2O Mg(OH)Cl ⎯→ MgO + HCl Hence, Anh. MgCl2 cannot be prepared by heating this hydrate. Because of this formation of HCl. Sea water cannot be used in marine boilers which corrodes the iron body. Anhydrous MgCl2 can be prepared by heating a double salt like. MgCl2.NH4Cl.6H2O as follows: – H2O strong ⎯→ MgCl2. NH4Cl ⎯⎯ ⎯→ MgCl2 + NH3 + HCl MgCl2 . NH4Cl . 6H2O ⎯⎯  (5) (6) (7)  It is a colourless crystalline solid, highly deliquescent and highly soluble in water. Sorel Cement is a mixture of MgO and MgCl2 (paste like) which set to hard mass on standing. This is used in dental filling, flooring etc. Anh. CaCl2 is used in drying gases and organic compounds but not NH 3 or alcohol due to the formation of CaCl2.8NH3 and CaCl2.4C2H5OH. s-Block Elements SULPHATES 1. Potassium Sulphate (K2SO4) It occurs in stassfurt potash beds as schonite K2SO4.MgSO4.6H2O and Kainite, KCl.MgSO4.3H2O from which it is obtained by solution in water and crystallisation. It separates from the solution as anh, crystals whereas Na2SO4 comes as decahydrate. Preparation: (1) It is prepared by the reaction of potassium chloride or hydroxide with concentrated. H 2SO4. 2KCl + H2SO4 ⎯→ K2SO4 + 2HCl ; 2KOH + H2SO4 ⎯→ K2SO4 + 2H2O (2) K2SO4.MgSO4.6H2O + 2KCl ⎯→ 2K2SO4 + MgCl2+ 6H2O Uses : It is used to prepare alum. It is a white crystalline solid and soluble in water. It is used as a fertilizer for tobacco and wheat. KCl + H2 SO4 2. (1) (2) (3) (4)  Reactions Charts Magnesium Sulphate (MgSO4): It occurs in nature as minerals kiesserite (MgSO 4.H2O), epsom salt (MgSO4.7H2O) and kainite (KCl.MgSO4.3H2O). Preparation: It is obtained by dissolving kieserite. MgSO4.H2O in boiling water and then crystallising the solution as a hepta hydrate. i.e. MgSO4.7H2O. It is called as Epsom salt. It is also obtained by dissolving magnesite in hot dil. H 2SO4. MgCO3 + H2SO4 ⎯→ MgSO4 + H2O + CO2 By dissolving dolomite (CaCO3.MgCO3) in hot dil. H2SO4 and removing the insoluble CaSO4 by filtration. CaCO3.MgCO3 (dolomite) + 2H2SO4 ⎯→ MgSO4 + CaSO4 + 2CO2 + 2H2O It is isomorphous with FeSO4.7H2O, ZnSO4.7H2O. s-Block Elements (1) Chemical Properties: Heating effect: When heated to 150°C, it changes to monohydrate. On further heating, it becomes anhydrous at 200°C. On strong heating, it decomposes into MgO. MgSO4.7H2O 150C ⎯⎯⎯⎯→ MgSO4.H2O 200C ⎯⎯⎯⎯→ MgSO4 strong ⎯⎯⎯⎯→ MgO + SO2 + O2. heating (2) Magnesium sulphate when heated with lamp black at 800°C produces SO2 and CO2 gases. 2MgSO4 + C ⎯→ 2MgO + 2SO2 + CO2 (3) It forms double salts with alkali metal sulphates, e.g., K 2SO4.MgSO4.6H2O. 3. Calcium Sulphate (Plaster of paris) CaSO4.½ H2O It occurs as anhydrite CaSO4, hemihydrate CaSO4.½H2O and as the dihydrate (CaSO4.2H2O) gypsum, alabaster or satin-spar. Preparation: (1) It is a hemihydrate of calcium sulphate. It is obtained when gypsum, CaSO 4·2H2O, is heated to 393 K. 393 K 2(CaSO4.2H2O) ⎯⎯ ⎯ ⎯→ 2(CaSO4).H2O + 3H2O Above 393 K, no water of crystallisation is left and anhydrous calcium sulphate, CaSO4 is formed. This is known as ‘dead burnt plaster’. It has a remarkable property of setting with water. On mixing with an adequate quantity of water it forms a plastic mass that gets into a hard solid in 5 to 15 minutes. (2) It can be prepared by reacting any calcium salt with either sulphuric acid or a soluble sulphate. CaCl2 + H2SO4 ⎯→ CaSO4 + 2HCl ; CaCl2 + Na2SO4 ⎯→ CaSO4 + 2NaCl Properties: It is a white crystalline solid. It is sparingly soluble in water and solubility decreases as temperature increases. It dissolves in dilute acids. It also dissolves in ammonium sulphate due to the formation of double sulphate, (NH4)2SO4.CaSO4.H2O. The setting process is exothermic. The process of setting takes place in stages. In the first stage, there is conversion of Plaster of Pairs into orthorhombic form of gypsum (setting step) and in the second stage orthorhombic form changes into monoclinic form (hardening step). The setting of Plaster of Paris may be catalysed by sodium chloride while it is retarded by borax or alum. Addition of alum to Plaster of Paris makes the setting very hard. The mixture is known as Keene’s cement. Dead plaster has no setting property as it takes up water only very slowly. A suspension of gypsum when saturated with ammonia and carbon dioxide forms ammonium sulphate, a nitrogenous fertilizer. 2NH3 + CaSO4 + CO2 + H2O ⎯→ (NH4)2 SO4 + CaCO3 When strongly heated with carbon, it forms calcium sulphide. CaSO4 + 4C ⎯→ CaS + 4CO Uses: For preparing blackboard chalk. In anhydrous form as drying agent. s-Block Elements Fertilizer 1. 2. Cynamide: It is an organic compound with the formula CN2H2. This white solid is widely used in agriculture and the production of pharmaceuticals and other organic compounds. Cyanamide is produced by hydrolysis of calcium cyanamide, which in turn is prepared from calcium carbide via the frank-Caro process. CaC2 + N2 ⎯→ CaCN2 + C ; frank-Caro process CaCN2 + H2O + CO2 ⎯→ CaCO3 + H2NCN (Cynamide) The main reaction exhibited by cyanamide involves additions of compounds containing an acidic proton. Water, hydrogen sulfide, and hydrogen selenide react with cyanmide to give urea, thiourea, and selenourea, respectively : H2NCN + H2E → H2NC (E) NH2 ; (E= O, S, Se) Fluorapatite: It is a phosphate mineral with the formula Ca5 (PO4)3. Cement Cement is a product obtained by combining a material rich in lime, CaO with other material such as clay which contains silica, SiO2 along with the oxides of aluminium, iron and magnesium. The raw materials for the manufacture of cement are limestone and clay. When clay and lime are strongly heated together they fuse and react to form cement clinker. This clinker is mixed with 2-3% by weight of gypsum (CaSO4.2H2O) to form cement. Thus important ingredients present in Portland cement are dicalcium silicate (Ca2SiO4) 26%, tricalcium silicate (Ca3SiO=5) 51% and tricalcium aluminate (Ca3Al2O6) 11%. Setting of cement : When mixed with water, the setting of cement takes place to give a hard mass. This is due to the hydration of the molecules of the constituents and their rearrangement. The purpose of adding gypsum is only to slow down the process of setting of the cement so that it gets sufficiently hardened. Uses : Cement has become a commodity of national necessity for any country next to iron and steel. It is used in concrete and reinforced concrete, in plastering and in the construction of bridges, dams and buildings. Common Names The names marked with asterisk (*) should be reference. You need not memorize them. Metal Ore name Lithium (Li) Spodumene Lepidolite Petalite Sodium (Na) *Washing soda *Baking soda *Sodium carbonate (soda ash/ washing soda) *Sodium chloride (rock salt or halite) *Sodium nitrate (Chile saltpeter) Salt cake Fusion mixture Sodium sesquicarbonate (trona) *Microcosmic salt Soda feldspar or sodium feldspar (albite) Potash feldspars or orthoclase or microcline or Potassium feldspars *Hypo *Sodium aluminium fluoride (cryolite) *Borax (Tincal) *Sodium sulphate (glauber’s salt) memorized with formulae. Others are given only for Formula LiAl(SiO3)2 KLi2Al(Al,Si)3O10(F,OH)2 LiAl(Si2O5)2 Na2CO3.10H2O NaHCO3 Na2CO3 NaCl NaNO3 Na2SO4 Na2CO3 + K2CO3 (eq. molar mix.) Na2CO3.NaHCO3.2H2O ( it is a double salt ) Na(NH4)HPO4.4H2O (it is obtained by mixing solutions of sodium phosphate and ammonium phosphate or chloride) Na2O. Al2O3. 6SiO2 K2O. Al2O3.6SiO2 Na2S2O3.5H2O Na3AlF6 Na2B4O7.10H2O Na2SO4.10H2O (Sodium sulfate is the sodium salt of sulfuric s-Block Elements acid. When anhydrous, it is a white crystalline solid of formula Na2SO4 known as the mineral thenardite; the decahydrate Na2SO4·10H2O is known as Glauber's salt) Potassium (K) Sodium aluminium silicate (Soda Feldspar) Sylvite Schonite Kainite *Carnallite *Indian saltpetre (Nitre) Beryllium (Be) Magnesium (Mg) Calcium (Ca) Strontium(Sr) Pearl ash Schonite Langbeinite Polyhalite *Potassium Alum Alunite or Alumstone Mica Feldspar Beryl Chrysoberyl Phenacite Bromalite *Baryta NaAlSi3O8 KCl K2SO4.MgSO4.6H2O MgSO4.KCl.3H2O MgCl2.KCl.6H2O KNO3 (used especially as a fertilizer and explosive) K2CO3 K2SO4.MgSO4.6H2O( it is a double salt) K2SO4.2MgSO4 K2SO4.MgSO4.2CaSO4.2H2O K2SO4. Al2 (SO4)3. 24H2O K2SO4. Al2 (SO4)3. 4Al(OH)3 K2O. 3Al2O3. 6SiO2.2H2O KAlSi3O8(K2O.Al2O3.6SiO2) 3BeO. Al2O3 6SiO2 BeO.Al2O3 BeSiO4 BeO Ba(OH)2 *Magnesite MgCO3 *Dolomite *Epsom salt Kieserite Asbestos Talc Brucite *Magnesia Artinite *Sorel cement (magnesia cement) *Quick lime *Slaked lime *Hydrolith *Calcium cynamide *Limestone (Marble / Whiting) Anhydrite *Gypsum *Fluorspar or Fluorite Phosphorite *Fluorapatite *Plaster of paris *Bleaching powder *Rock phosphate Wollastonite Colmanite Strontianite Celestite Barytes or Heavy spar MgCO3.CaCO3 MgSO4.7H2O MgSO4.H2O CaMg3 (SiO3)4 Mg(Si2O5)2 Mg (OH)2 Mg(OH)2 MgO MgCO3.Mg(OH)2 .3H2O Mg4Cl2(OH)6(H2O)8 CaO Ca(OH)2 CaH2 CaCN2 OR CaNCN CaCO3 CaSO4 CaSO4.2H2O CaF2 Ca3 (PO4)2 3Ca3 (PO4)2.CaF2 OR Ca5(PO4)3F CaSO2.½H2O CaOCl2 Ca3 (PO4) 2 CaSiO2 2CaO.3Ba2O3.5H2O SrCO3 SrSO4 BaSO4 s-Block Elements Periodic Properties of s-Block Properties Order Thermal stability LiH > NaH > KH > RbH > CsH Basic strength