2. 1 S-BLOCK ELEMENTS The s-Block Element

3. 2 Group I Elements (Alkali Metals) They have similar chemical properties. They are soft metals with fixed O.N. = +1 in their compounds. Electronic configuration Atomic radius (Å) Ionic radius I.E.M.P.DensityE.N.  H hyd of cation LiHe 2s 1 1.230.605201810.541.0-519 NaNe 3s 1 1.570.95496980.970.9-406 KAr 4s 1 2.031.33419630.860.8-322 Rb Kr 5s 1 2.161.48403391.530.8-293 CsXe 6s 1 2.351.69375291.870.7-264 FrRn 7s 1 --1.76--27--0.7--

4. 3 Group II Elements (Alkaline Earth Metals) Mg and Ca are the most abundant elements and Ra is the most scarce element which is unstable and radioactive. Electronic configuration Atomic radius ( Å) Ionic radius I.E.M.P.DensityE.N.  H hyd of cation 1st2nd BeHe 2s 2 0.900.31900176012771.851.5-2494 MgNe 3s 2 1.360.6573714506501.741.2-1921 CaAr 4s 2 1.740.9959011508381.551.0-1577 SrKr 5s 2 1.911.1354810607682.601.0-1443 BaXe 6s 2 1.981.355029657143.500.9-1305 RaRn 7s 2 --1.405.99797005.000.9--

5. 4 Atomic and ionic radius Atomic radius and ionic radius increase An addition of one more shell; ENC decrease

6. 5  1 Compare the atomic radius of alkali and alkaline earth metals Alkali metals > alkaline earth metals because Increase in nuclear charge > increase in shielding effect ENC increases and electron experience a larger nuclear attraction.

7. 6  2 Compare the atomic and ionic radius Ionic radius < corresponding atomic radius Same nuclear charge, weaker shielding effect ENC increases, stronger nuclear attraction towards electrons.

8. 7 Melting and boiling point M.P. decreased ENC decreased so that nuclear attraction towards electrons ( metallic bond strength)decreased.

9. 8  4. Difference between m.p. and b.p. of s-block metals b.p. > m.p. a lot Most of the metallic bonds remains in the liquid state; nearly all bonds in liquid state have to be broken on vaporization.

10. 9 Atomic volume Atomic volume increased Atomic size increases, Metallic bond strength decreases.

11. 10  5. Compare the atomic volumes of s-block metals Alkali metals > Alkaline earth metals Atomic size of alkaline earth metals < alkali metals Metallic bond strength of alkaline earth metals > alkali metals

12. 11 Density Density increased Atomic mass increased to a greater extent than atomic volume.

13. 12 Alkali metals < Alkaline earth metals Atomic mass of alkaline earth metals > alkali metals Atomic volume of alkaline earth metals < alkali metals Density = atomic mass / atomic volume  6. Compare the density of s-block metals

14. 13 Ionization energy I.E. decreased ENC decreases due to addition of shells. I.E. of alkaline earth metals > alkali metals ENC increases

15. 14 Electron being removed is from the inner shell; the electron thus experience a larger nuclear attraction. Besides, ENC of an ion would be much greater than the corresponding atom as the shielding effect is weaker.  7. Difference between 1 st and 2 nd I.P. of alkali metals

16. 15 Summary On passing down the group: 1.Atomic radius and ionic radius increased An addition of one more shell; ENC decrease 2.I.E. decreased ENC decreased due to addition of shells. 3.M.P. decreased ENC decreased so that nuclear attraction towards electrons decreased. 4.Density increased atomic mass increased to a greater extent than atomic volume. 5.E.A. decreased ENC decreased so that tendency to accept e - decreased. 6.Reducing power and reactivity increased I.E decrease and reduction potential become more negative. 7.Enthalpy of hydration of cation less negative Electrostatic interaction between the polar water molecules and ions become less as the ionic radius increases.

17. 16 Summary of physical properties of Group I and IIA elements m.p. and b.p. ; density and I.P. Atomic and ionic radii ; atomic volume m.p. and b.p. ; density and I.P.

18. 17 Variaiton in chemical properties Owing to the low value of 1 st I.P., alkali metals are relatively more easily to form X +, and the resulting compound is quite stable. The sum of 1 st and 2 nd I.P. of alkaline earth metals is not too low, yet the lattice energy recorded on forming the ionic compounds is large enough for the formation of X 2+.

19. 18 Try to account for the following redox potentials: E o / V Li(s)  Li + (aq) + e – 3.04V Cs(s)  Cs + (aq) + e – 2.93V Rb(s)  Rb + (aq) + e – 2.93V K(s)  K + (aq) + e – 2.92V Na(s)  Na + (aq) + e – 2.71V Reducing power and reactivity of s-block elements

20. 19 Reducing power and reactivity of s-block elements The redox potential: M(s)  M + (aq) + e depends on: 1. the formation of separate atom from crystal lattice M(s)  M(g)  H sub = heat of sublimation (+ve) 2. the formation of gaseous ion from gaseous atom M(g)  M + (g) + e I.E. ionization energy (+ve) 3. the formation of hydrated ion from gaseous ion M + (g) + aq  M + (aq)  H hyd = hydration energy (-ve)

21. 20 The overall enthalpy change (  H) =  H sub + I.E. +  H hyd  H more -ve  the greater the redox potential i.e. the stronger the reducing agent. Reducing power and reactivity of s-block elements

22. 21 From Na to Cs, the reduction potential increased. But Li has greatest reduction potential. On passing down the group, both  H atm, I.E. decrease but  H hyd also become less -ve. But  H sub and I.E. decrease to a greater extent than the  H hyd,  H(overall) is more negative. Reducing power and reactivity of s-block elements

23. 22 Li is an exceptional case, it has the greatest redox potential. It is because the size of Li + is very very small (it belongs to 2nd period ),  H hyd is exceptionally more -ve. Therefore  H(overall) is thus more negative. Reducing power and reactivity of s-block elements

24. 23 Variation in chemical properties Reactions: 1.With air - All tarnish in air (that is, forming a film of oxide on the surface), therefore they are stored in paraffin oil. When burnt in sufficient amount of oxygen : Kind of oxides Elements which form this type of oxide in adequate supply of air normal oxide O 2- Li, Mg, Ca, Sr peroxidesNa, Ba superoxide O 2- K, Rb, Cs

25. 24 Reaction with air Dot and cross diagram for oxide O 2- ion and peroxide O 2 2- ion : O – O bond can be easily broken OO        2- Size of peroxide ion > size of oxide ion

26. 25 Reaction with air Li + ion is extremely small, it is not possible for sufficient number of peroxide ions to surround the Li + ion with causing repulsion between the anions, therefore only normal oxide exists. The larger peroxide and superoxide anions are stabilized by larger cations due to limiting radius ratio.

27. 26 Stability of oxide ion > peroxide and superoxide ion Ba 2+, being the largest ion, has the weakest polarizing power ; electron cloud of the peroxide ion will be distorted by other group IIA [AND also Li + ] metal ions and become unstable. Ba 2+ is the biggest ion in Group IIA [slightly larger than K + ], no severe repulsion would occur between these large peroxide ion when surrounding Ba 2+ in the lattice.  9. Why Group IIA elements form normal oxides, except barium ?

28. 27 Reaction with air In case of Li, Mg, Ca, Sr and Ba, the final products will be a mixture of nitrides, carbonates together with the oxides. Only Li in group IA would form Li 3 N (lithium nitride). N 3- ion is hard to form, why ? 6 Li + N 2  2 Li 3 N Li 3 N + 3H 2 O  3LiOH + NH 3

29. 28 With water - All (except Be) reacts to give out hydrogen. 2H 2 O + 2e -  2OH - + H 2 at pH = 7 E = -0.41V 2M + 2H 2 O  2MOH + 2H 2  E = +ve (spontaneous) But the vigor of reaction: K > Na > Li although  E of Li is greatest. Why?  E shows the equilibrium position (i.e. the reaction is spontaneous or not ),  E increased implies that equilibrium lies on the product side. Rate of reaction must consider the E act (activation energy). From the information given, the rate of Li is the slowest among the three. That is E act for Li is the highest, so the rate is relatively slow. Reaction with water

30. 29 The rate of reaction increases on passing down the group. Reactivity of alkali metals towards water is much higher than alkaline earth metals. Magnesium reacts with hot water and steam to give magnesium hydroxide and magnesium oxide respectively. Reaction with water

31. 30 All react vigorously and explosively. 2M + 2H +  H 2 + 2M + But reactions between sulphuric acid and Ca, Sr, Ba become less vigorous after the reaction starts due to the formation of insoluble layer of sulphates. Ca + H 2 SO 4 CaSO 4 (s) + H 2 Reaction with acid

32. 31 With non-metal - All combine with X 2, S and O 2, P or even H 2 at suitable temperature. Li, Mg, Ca, Sr and Ba also combine directly with nitrogen. Ca + H 2 CaH 2 calcium hydride 2 K + SK 2 S potassium sulphide 3 Mg + N 2 Mg 3 N 2 magnesium nitride Reaction with non-metals

33. 32 Oxides All are white crystalline solid, ionic and strongly basic in character. They are hydrolysed by water to form corresponding hydroxides. Degree of hydrolysis increases down the group, since oxide become more ionic. O 2- + H 2 O2 OH - BeO, being exceptional case, is amphoteric : BeO + 2HCl  BeCl 2 + H 2 O BeO + 2OH - + H 2 O  Be(OH) 4 2- beryllate

34. 33 Na 2 O is more basic. Na + has a weaker polarizing power than Mg 2+ (as the latter one has a higher charge/radius ratio OR charge density), electron in O 2- ion is more available to attack hydrogen in water molecule. More hydroxide ion is thus formed.  10. Compare the basic strength of Na 2 O and MgO

35. 34 Hydrides Formed by heating the element in hydrogen gas (at 400 ℃ or above) ; Strong reducing agents ; Hydrolysed by water to form hydrogen gas and solution or suspension of hydroxides ; Readiness of hydrolysis increases down the group since the hydride is more ionic ; reaction is more vigorous for alkali metals than for alkaline earth metals

36. 35 Chlorides All are white crystalline solid, soluble in water to form hydrated ion. NaCl(s) + aq. Na + (aq) + Cl - (aq) Hydrated sodium and chloride ion No hydrolysis and thus a neutral solution But for MgCl 2, which is partially ionic, hydrolysed by water to give a slightly acidic solution. MgCl 2 (s) + 6 H 2 O Mg(H 2 O) 6 2+ + 2 Cl - (aq) Mg(H 2 O) 6 2+ + H 2 O Mg(H 2 O) 5 (OH) + + H 3 O + Be(H 2 O) 4 2+ + H 2 O Be(H 2 O) 3 (OH) + + H 3 O +

37. 36 Thermal stability of other compounds For a large polarizable anion (e.g. HCO 3 -, CO 3 2-,NO 3 -, SO 4 2- ), the stability depends on the polarizing power of the cation. If the cation can distort the electron cloud of the anion so much that the bonds (e.g. C-O bond in carbonate) is weakened, the bond will be easily broken on heating to give metallic oxides and gas(es) (CO 2 for carbonate).

38. 37 Thermal stability  MgCO 3 MgO + CO 2  MgSO 4 MgO + SO 3  2 Mg(NO 3 ) 2 2 MgO + 2 NO 2 + O 2  2 NaNO 3 2 NaNO 2 + O 2

39. 38 Na 2 CO 3 is thermally stable because polarizing power of Na + is weaker than Mg 2+ (as the latter one has a higher charge/radius ratio / charge density), electron cloud of the carbonate ion is much distorted by Mg 2+ that the C – O bond is weakened and thus more easily broken when heated  11. Compare the stability of Na 2 CO 3 and MgCO 3

40. 39 Polarising power of cation decreases on passing down the group as the size of the cation become larger. Most group I salts are thermally stable except for those of lithium. While group II salts are relatively less stable to heat. ( Note that only lithium carbonate is thermally unstable among group I carbonates.). Thermal stability

41. 40 Thermal stability Some sodium and potassium salts are decomposed when heated : 2 NaNO 3 2 NaNO 2 + O 2 2 NaHCO 3 Na 2 CO 3 + CO 2 + H 2 O

42. 41 Solubility of salts in water All group I compounds are practically soluble.The solubility increases as heat of hydration is more negative than lattice energy.  H hyd  H latt  H hyd -  H latt LiI-824-763-61 NaI-711-703-8 KI-627-647+20 RbI-598-624+26 CsI-569-601+32 LiF-1034-1039+6 NaF-927-919-2 KF-837-817-20 RbF-808-779-29 CsF-779-730-49

43. 42 Lattice energy depends on the sum of the ionic radii while the hydration energy depend on ionic radius of the individual ions ; both would decrease as size of ions increases. Hydration energy of a compound is contributed by both the cation and anion. Solubility of salts in water

44. 43 Hydration energy contributed by anion is small, i.e. the hydration energy mainly contributed by the cation. On passing down the group, hydration energy decrease greatly / tremendously. The sum of ionic radii only increase slightly as the size of anion is large, the decrease in lattice energy is small.  H soln =  H hyd  – LE   H soln become less negative (more positive) on passing down the group. Solubility of salts with large anion (e.g. I -, SO 4 2-, CO 3 2- ) in water

45. 44 Solubility of salts with large anion (e.g. I -, SO 4 2-, CO 3 2- ) in water CompoundsSolubility / x 10 -3 mol dm -3 Described as MgSO 4 3600soluble CaSO 4 11sparingly soluble SrSO 4 0.62insoluble BaSO 4 0.009insoluble

46. 45 Lattice energy decrease more rapidly than the hydration energy on passing down the group, hydration energy mainly depends on the small anion and would not change much; decrease in lattice energy mainly determines the solubility of the salt.  H soln =  H hyd  – LE   H soln become more negative (less positive) on passing down the group. Solubility of salts with small anion (e.g. F -, OH - ) in water

47. 46 Solubility of salts with small anion (e.g. F -, OH - ) in water CompoundsSolubility / x 10 -3 mol dm -3 Described as Mg(OH) 2 0.020insoluble Ca(OH) 2 1.5slightly soluble Sr(OH) 2 3.4soluble Ba(OH) 2 15soluble

48. 47 Compare the solubility of salts of alkali and alkaline earth metals An increase in charge will increase lattice energy to a greater extent than hydration energy, salts of alkaline earth metals are generally less soluble than that of alkali metals, and doubly charged anions give more insoluble compounds.

49. 48 General characteristics of s-block elements Fixed oxidation state The only possible positive oxidation state shown by the elements is equal to the total number of electron in the outermost shell. This oxidation state corresponds to the loss of sufficient number of electrons to achieve the octet configuration ns 2 np 6, thus only forms compounds in which they obtain the octet configuration. The loss of more than the valence electron requires too much ionization energy, thus prevents these metals from showing an oxidation number other than the one equal to their group number.

50. 49 Ability to form complexes Owing to the lack of underlying (inner) low energy vacant orbital, s-block elements rarely form complexes. The cations which form stable complexes normally carrying a high charge / radius ratio, resulting in larger electrostatic attraction between the central ions and the ligands. General characteristics of s-block elements

51. 50 Ability to form complexes Group I metal ions cannot form hydrated ions of definite formula in aqueous solution, though they can by hydrated to certain extent. Lithium ion, which has the smallest size, show certain degree of hydration in crystal of its salts.

52. 51 Group II metals ions has higher charge/radius ratio and they have higher tendency to form complexes. Beryllium forms many complex but barium forms very few. e.g. BeF 3 - BeF 4 2- Be(H 2 O) 4 2+ The most important complex for Mg is chlorophyll, which has a very complicated structure with fused rings; the Mg atom being at the center of the rings bonded to 4 nitrogen atoms. Ability to form complexes

53. 52 Ca 2+ and Mg 2+ form stable complex with strong complexing agents. e.g. ethylenediaminetetraacetic acid ( EDTA ) which has 4 functional oxygen atoms and 2 donor N- atoms per molecule. More discussion on d-block elements. Ability to form complexes

54. 53 Abnormality of lithium and its compounds among group IA 1 Lithium carbonate and hydroxide are decomposed by heat. 2 Lithium carbonate, hydroxide and fluoride are insoluble in water. 3 Lithium forms only normal oxide when reacting with oxygen. 4 Lithium forms nitride when heated in air. 5 Lithium ion is highly hydrated in water, resulting in lowest mobility. 6 Lithium hydroxide is not a strong base. 7 Almost all lithium salts are hydrated in its crystal lattice.

55. 54 Reason : Exceptional small size of Li + ion, e.g. forming nitride : 6 Li(s) + N 2 (g)2 Li 3 N(s) 6 Li(g) + 2 N(g)2 N 3- (g) + 6 Li + (g) Abnormality of lithium and its compounds among group IA The highly negative LE of Li 3 N offset the energy required to ionize the nitrogen gas to nitride ion.

56. 55 Diagonal relationship between magnesium and lithium 1.Both only form normal oxide. 2.Both give nitrides when heating in air. 3. Carbonates,sulphates,hydroxides are decomposed by heat to metallic oxides. 4.Carbonates, hydroxides are insoluble in water. Reasons: Effective nuclear charge increases on passing along the period but decreases on passing down the group, so Li + and Mg 2+ have similar effective nuclear charge which in turn affecting its polarizing power.

57. 56 Flame Test for Metal Ions: ElementFlame ColourElementFlame Colour LithiumredCalciumbrick red Sodiumgolden yellowStrontiumblood red PotassiumlilacBariumapple green Rubidiumviolet Caesiumviolet