1. 40
The s-Block Elements
40.1 Characteristic Properties of the s-Block Elements
40.2 Variation in Properties of the s-Block Elements
40.3 Variation in Properties of the Compounds of the s-Block Elements

2. The Syllabus 8.1 Characteristic Properties Metallic character
Low electronegativity
Formation of basic oxides and hydroxides
Fixed Oxidation state in their compounds
Weak tendency to form complexes
Flame colours of salts – flame test

3. The Syllabus
8.2 Variation in properties of the s-block elements and their compounds
Variations in atomic radii, ionisation enthalpies, hydration enthalpies and melting points.
Interpretation of these variations in terms of structure and bonding.
Reactions of the elements with oxygen and water. Reactions of the oxides with water, dilute acids and dilute alkalis.
Relative thermal stability of the carbonates and hydroxides. Relative solubility of the sulphates(VI) and hydroxides

4. Notes p. 1
s-Block elements:
Consists of Group IA and Group IIA elements
Outermost electron shell: ns1 ns2
Highly reactive metals
Good reducing agents
Fixed oxidation states +1 for Group I elements +2 for Group II elements

5. Characteristic Properties of the s-Block Elements
40.1
Characteristic Properties of the s-Block Elements

6. Metallic Character (not mentioned in notes)
40.1 Characteristic Properties of the s-Block Elements (SB p.38)
Metallic Character (not mentioned in notes)
Group I elements:
Silvery in colour, tarnish rapidly in air
∴ keep immersed under paraffin oil or in vacuum sealed tubes
Soft, low boiling and melting points
∵ weak metallic bond due to only 1 e– is contributed to form bonds
Low density
∵ body-centred cubic structure -- have more spaces
Cutting Rubidium

7. Group I elements:
Lithium
Sodium
Potassium
Caesium
Rubidium

8. 40.1 Characteristic Properties of the s-Block Elements (SB p.39)
Some information about Group I elements
Group I metal
Atomic radius (nm)
Ionic radius (nm)
Crystal structure
Melting point (C)
Boiling point (C)
Density (g cm–3)
Abundance on earth (%) Li Na K Rb Cs Fr
0.152
0.186
0.231
0.244
0.262
0.270
0.060
0.095
0.133
0.148
0.169
0.176 b —
180.5
97.8
63.7
39.1
28.4 27
1330
890
774
688
690
680
0.53
0.97
0.86
1.53
1.87
0.0020
2.36
2.09
Trace
“b” denotes body-centred cubic structure

9. harder and higher boiling and melting points than Group I counterparts
40.1 Characteristic Properties of the s-Block Elements (SB p.39)
Group II elements:
silvery in colour
harder and higher boiling and melting points than Group I counterparts
∵ stronger metallic bond due to 2e– are contributed to form bond and smaller atomic sizes
show different crystal structures

10. Group II elements:
Calcium
Beryllium
Magnesium
Radium
Strontium
Barium

11. 40.1 Characteristic Properties of the s-Block Elements (SB p.39)
Some information about Group II elements
Group II metal
Atomic radius (nm)
Ionic radius (nm)
Crystal structure
Melting point (C)
Boiling point (C)
Density (g cm–3)
Abundance on earth (%) Be Mg Ca Sr Ba Ra
0.112
0.160
0.197
0.215
0.217
0.220
0.031
0.065
0.099
0.113
0.135
0.140 h f b —
1278
648.8
839
769
729
697
2477
1100
1480
1380
1640
1140
1.85
1.75
1.55
2.54
3.60
5.0
2.33
4.15
0.038
0.042
Trace
“h”, “f” and “b” denote hexagonal close-packed, face-centred cubic and body-centred cubic structures respectively

12. Variation in Physical Properties
Atomic Radius and Ionic Radius (notes p. 1)

13. Question: The atomic and ionic radii increase down the Groups, why?
41.3 Variation in Properties of the s-Block Elements (SB p.52)
Question:
The atomic and ionic radii increase down the Groups, why?
∵ outermost shell electrons become further away, and more inner shells shielding the outermost shell electrons
 attraction between the nucleus and the outermost shell electrons decreases
 atomic and ionic radii increase

14. Question:
Atomic and ionic radii decrease when going from Group I to II in each period, why?
∵ Group II elements have 1 more proton and electron than Group I elements. Increase in nuclear charge outweighs the increase in shielding effect of additional electron of the same shell.
 atomic and ionic radii decrease

15. 41.3 Variation in Properties of the s-Block Elements (SB p.52)
Question:
Ionic radius of any Group I or II element is smaller than the atomic radius, why?
∵ after losing the outermost shell electron(s), there is one electron shell less in the cation than in the atom.
Increase in p/e ratio

16. Ionization Enthalpy (notes p. 2)
41.3 Variation in Properties of the s-Block Elements (SB p.53)
Ionization Enthalpy (notes p. 2)

17. 41.3 Variation in Properties of the s-Block Elements (SB p.54)
Variations in the 1st, 2nd and 3rd ionization enthalpies of Group II elements

18. 1st I.E. is much smaller than 2nd I.E. for Gp. I elements
For the 1st I.E., electron is further away from the nucleus and shielding effect of inner shell electrons  small 1st I.E.
For 2nd I.E., electron is removed from stable noble gas configuration and higher effective nuclear charge  large 2nd I.E.

19. The ionization enthalpies decrease down the Groups
Reason:
atomic sizes increase down the group
 the outermost shell electron(s) is/are further away from the nucleus, they will be better shielded by inner electron shells.
 less attractive force experienced
 less energy is required to remove the electrons
Because of the high I.E., Li and Be forms a few covalent compounds instead of forming Li+ and Be2+ respectively.

20. Low Electronegativity
40.1 Characteristic Properties of the s-Block Elements (SB p.41, notes p. 3))
Low Electronegativity
All have low electronegativity values
∵ the outermost electron shell is effectively shielded by inner electron shells.
- Low effective nuclear charge.
Decrease when going down the group
∵ the outermost electron shell are further away from nucleus
- increase in shielding effect.

21. 40.1 Characteristic Properties of the s-Block Elements (SB p.41)
Group II elements are relatively more electronegative than Group I counterparts
∵ higher nuclear charge, stronger attraction to outermost shell electrons
Group I element
Electro-negativity
Group II element Li Na K Rb Cs Fr
1.0
0.9
0.8
0.7 — Be Mg Ca Sr Ba Ra
1.5
1.2

22. Characteristic Flame Colours of Salts
40.1 Characteristic Properties of the s-Block Elements (SB p.43)
Characteristic Flame Colours of Salts
The outermost shell electrons of Group I & II elements are weakly held
 The electrons can be excited to higher energy levels on heating
 When electrons return to ground state, radiations are emitted
 The radiations fall into the visible light region
 The flame colour is a characteristic property of the element

23. Flame Test

24. 40.1 Characteristic Properties of the s-Block Elements (SB p.43)

25. Flame colours

26. Weak tendency to form complexes (not mentioned in notes)
40.1 Characteristic Properties of the s-Block Elements (SB p.43)
Weak tendency to form complexes (not mentioned in notes)
Complex: Polyatomic ion or neutral molecule formed when molecular or ionic gropups (called ligands) form dative covalent bonds with a central ion.
Group I & II elements seldom form complex:
s-block ions do not have low energy vacant orbitals available for dative covalent bonds.
Low ionic charge

27. Melting Point (notes p. 4)
41.3 Variation in Properties of the s-Block Elements (SB p.55)
Melting Point (notes p. 4)
Variations in melting points of Groups I and II elements

28. Observations: • melting point decreases as going down Groups I and II
Reason:
the ionic size of the elements increases
 attraction between ions and electrons becomes weaker
 metallic bond is weaker

29. Observations:
• melting points of Group II elements are much higher than those of Group I elements
Reason:
no. of valence electrons per mole contributed to the delocalized electron sea is greater.
Group II elements have higher ionic charge
 the attractive force between ions and electrons are stronger
 metallic bond is stronger

30. 41.3 Variation in Properties of the s-Block Elements (SB p.56)
Observations:
• irregularity in the general decrease in melting point down Group II elements
Reason:
different metallic crystal structures of the Group II elements
Group II metal
Crystal structure Be Mg Ca Sr Ba Ra h f b —

31. Extraction of sodium (not in syllabus)
Downs Cell

32. Manufacture of sodium hydroxide
graphite anodes
chlorine
used brine
mercury alloyed
with sodium
flow of mercury
flowing mercury (as cathode)
saturated brine + -
Mercury (recycle)
Water
Flowing mercury cell

33. During electrolysis, chlorine is liberated at the anode and sodium at the cathode.
At anode (graphite): 2Cl-(aq)  Cl2(g) + 2e-
At cathode (mercury): Na+(aq) + e-  Na(s);
Na(s) + Hg(l)  Na/Hg(l)
sodium amalgam

34. Flowing mercury cell
Q. 1b; Q.8

35. Hhyd must be negative value.
40.3 Variation in Properties of the s-Block Elements (SB p.56, notes p. 8)
Hydration Enthalpy
Xn+(g) + aq  Xn+(aq)
Hydration enthalpy (Hhyd) is the amount of energy released when one mole of aqueous ions is formed from its gaseous ions.
Hhyd must be negative value.
Hhyd depends on charge density charge/size
 Higher the charge, stronger the attraction, more energy released
 Smaller the size, stronger the attraction, more energy released

36. M+
Variations in hydration enthalpy of Groups I and II elements

37. •. magnitude of hydration enthalpies become smaller (less
• magnitude of hydration enthalpies become smaller (less negative) as going down the Groups
Reason:
the ionic size of the elements increases down the group, the charge density decreases
 the attractive force between water molecules and ions becomes weaker
 the hydration enthalpy becomes less negative Down the group, fewer molecules of water of crystallization
Na2CO3.10H2O MgSO4.7H2O MgCl2.6H2O
K2CO3.2H2O CaSO4.2H2O CaCl2.6H2O
SrSO4 BaCl2.2H2O

38. Observations:
• hydration enthalpies of Group II ions are more negative than those of Group I ions
Reason:
Group II ions have higher charge and smaller size
 charge density is much higher that of Group I ions
 the attractive force would be much stronger

39. Lattice Enthalpies of Group I Halides (p.10)

40. Lattice Enthalpies of Group I Halides (p.10)
Good agreement between calculated and measured value.
Why?
Lattice Enthalpies decrease down the group:
Reasons:
Size increase
Internuclear distance increase
Attractive force between opposite ions decrease

41. Lattice Enthalpies of Group II Halides (p.11)
Discrepancies occurred between calculated and measured values.
Reason:
Covalent characters occurred in small cations.
Group II Halides have a higher lattice enthalpies than Group I Halides.
Reason:
Higher charge; smaller size.

42. Formation of Hydroxides – reactions with water
40.1 Characteristic Properties of the s-Block Elements (SB p.43, notes p. 13)
Formation of Hydroxides – reactions with water
All Group I metals react with H2O to form metal hydroxides and H2 gas
e.g. 2Na(s) + 2H2O(l)  2NaOH(aq) + H2(g)
2K(s) + 2H2O(l)  2KOH(aq) + H2(g)
Li+H2O
Na +H2O

43. Rb+H2O
K+H2O
Cs+H2O

44. e.g. Ca(s) + 2H2O(l)  Ca(OH)2(aq) + H2(g)
40.1 Characteristic Properties of the s-Block Elements (SB p.43)
All Group II metals (except Be) react with H2O to form metal hydroxides and H2 gas (Mg reacts with hot water).
e.g. Ca(s) + 2H2O(l)  Ca(OH)2(aq) + H2(g)
Sr(s) + 2H2O(l)  Sr(OH)2(aq) + H2(g)
Be does not react with H2O(l or g)

45. Barium + water
Strontium + water

46. Formation of Basic Oxides
40.1 Characteristic Properties of the s-Block Elements (SB p.41, notes p. 14)
Formation of Basic Oxides
Group I elements
Produce more than one type of oxides (except Li)
All are ionic
Three types of oxides: normal oxides (monoxides), peroxides, superoxides
Relationship between three oxides:
O2–  O22–  2O2–
monoxide peroxide superoxide

47. Li forms the monoxide only 4Li(s) + O2(g)  2Li2O(s)
40.1 Characteristic Properties of the s-Block Elements (SB p.41)
Li forms the monoxide only
4Li(s) + O2(g)  2Li2O(s)
180C
Na forms the monoxide and peroxide when O2 is abundant
4Na(s) + O2(g)  2Na2O(s)
2Na2O(s) + O2(g)  2Na2O2(s)
180C
300C

48. K forms the monoxide, peroxide and superoxide
40.1 Characteristic Properties of the s-Block Elements (SB p.41)
K forms the monoxide, peroxide and superoxide
4K(s) + O2(g)  2K2O(s)
2K2O(s) + O2(g)  2K2O2(s)
K2O2(s) + O2(g)  2KO2(s)
180C
300C
3000C
Rb, Cs also forms superoxides
Rb2O2(s) + O2(g)  2RbO2(s)
Cs2O2(s) + O2(g)  2CsO2(s)
3000C

49. Group I element Monoxide Peroxide Superoxide
41.2 Characteristic Properties of the s-Block Elements (SB p.45)
Group I element
Monoxide
Peroxide
Superoxide Li Na K Rb Cs
Li2O
Na2O
K2O
Rb2O
Cs2O —
Na2O2
K2O2
Rb2O2
Cs2O2
KO2
RbO2
CsO2

50. Li does not form peroxides or superoxides Reason:  Li+ is small
40.1 Characteristic Properties of the s-Block Elements (SB p.42 notes p. 14)
Li does not form peroxides or superoxides
Reason:
 Li+ is small
 high polarizing power
 serious distortion on electron cloud of peroxide or superoxide (large polyatomic anions)
 more distortion , more unstable
 Li2O2 and LiO2 do not exist
K+, Rb+ and Cs+ ions are large
 Low polarizing power  peroxides and superoxides are relatively stable

51. Form normal oxides only, except Sr, Ba which can form peroxides.
41.2 Characteristic Properties of the s-Block Elements (SB p.46, notes p. 14)
Group II Elements
Form normal oxides only, except Sr, Ba which can form peroxides.
All are basic (except BeO which is amphoteric), why?
2Be(s) + O2(g)  2BeO(s)
2Mg(s) + O2(g)  2MgO(s)
2Ca(s) + O2(g)  2CaO(s)
2Ba(s) + O2(g)  2BaO(s)
2BaO(s) + O2(g) 2BaO2(s)

52. Strontium + air
Barium + air

53. Be, Mg, Ca peroxide do not exist, why?
40.1 Characteristic Properties of the s-Block Elements (SB p.43)
Group II
element
Normal oxide
Peroxide
Superoxide Be Mg Ca Sr Ba
BeO
MgO
CaO
SrO
BaO —
SrO2
BaO2
Be, Mg, Ca peroxide do not exist, why?
Reason:
 High charge density  high polarizing power
 serious distortion on electron cloud of the peroxide ion

54. Reactions of Oxides of s-Block Elements
40.3 Variation in Properties of the compounds of the s-Block Elements ( p.59)
notes p (e)
Dissolution of Na2O2 in H2O containing phenolphthalein
Reactions of Oxides of s-Block Elements
Reaction with Water
Group I oxides react with H2O to form hydroxides
Normal oxides:
e.g. Li2O(s) + H2O(l)  2LiOH(aq)
Peroxides:
e.g. Na2O2(s) + 2H2O(l)
 2NaOH(aq) + H2O2(aq)
Superoxides:
e.g. 2KO2(s) + 2H2O(l)  2KOH(aq) + H2O2(aq) + O2(g)

55. e.g. CaO(s) + H2O(l)  Ca(OH)2(aq) (weakly alkaline)
Group II oxides (except BeO, MgO) react with H2O to form a weakly alkaline solution
e.g. CaO(s) + H2O(l)  Ca(OH)2(aq) (weakly alkaline)
The basicity of all Group II oxides increases down the group
BeO is amphoteric
BeO(s) + 2H+(aq)  Be2+(aq) + H2O(l)
BeO(s) + 2OH–(aq) + H2O(l)  [Be(OH)4]2–(aq)
hot
MgO is slightly soluble in water, but dissolves in acids to form salts
BaO2(s) + 2H2O(l)  Ba(OH)2(aq) + H2O2(aq)

56. e.g. CaO(s) + 2HCl(aq)  CaCl2(aq) + H2O(l) Peroxides:
Variation in Properties of the compounds of the s-Block Elements (p.60, not mentioned in notes)
Reaction with Acids
All oxides of s-Block elements are basic except BeO which is amphoteric
Normal oxides:
e.g. CaO(s) + 2HCl(aq)  CaCl2(aq) + H2O(l)
Peroxides:
e.g. Na2O2(s) + 2HCl(aq)  2NaCl(aq) + H2O2(aq)
Superoxides:
e.g. 2KO2(s) + 2HCl(aq)  2KCl(aq) + H2O2(aq) + O2(g)

57. BeO is amphoteric, it reacts with NaOH to give Na2Be(OH)4
40.3 Variation in Properties of the compounds of the s-Block Elements (p.60)
Reaction with Alkalis
No reaction between the oxides of s-block elements with alkalis except BeO
BeO is amphoteric, it reacts with NaOH to give Na2Be(OH)4
BeO(s) + 2NaOH(aq) + H2O(l)  Na2Be(OH)4(aq)

58. Relative Thermal Stability of the Carbonates and Hydroxides
40.3 Variation in Properties of the compounds of the s-Block Elements (p.60)
notes p. 15, 18
Relative Thermal Stability of the Carbonates and Hydroxides
Thermal stability refers to the resistance of a compound to decomposition on heating
The higher the thermal stability of a compound, the higher is the temperature needed to decompose it
The thermal stability of ionic compounds depends on:
(1) charges &
(2) sizes of ions

59.  The compound tends to be less thermal stable
40.3 Variation in Properties of the compounds of the s-Block Elements (p.61)
notes p. 18
Compound with large polarizable polyatomic anion (large electron cloud, as shown in notes), the thermal stability depends on the polarizing power (charge density) of cations
 The stronger the polarizing power, the electron cloud of anion will be distorted to greater extent
 The compound tends to be less thermal stable

60. Group II carbonates/hydroxides are less stable than Group I
40.3 Variation in Properties of the compounds of the s-Block Elements (p.61)
Group II carbonates/hydroxides are less stable than Group I
Group II ions are smaller and have a higher charge than Group I ions in the same period
 Greater polarizing power
 The carbonates and hydroxides of Group II metals are less stable on heating
e.g. K2CO3 is stable upon heating while CaCO3 decomposes on heating

61. e.g. MgCO3(s)  MgO(s) + CO2(g) Ca(OH)2(s)  CaO(s) + H2O(g)
40.3 Variation in Properties of the compounds of the s-Block Elements (p.61)
Most carbonates and hydroxides of Group II metals readily undergo decomposition on heating to give oxides (more stable)
e.g. MgCO3(s)  MgO(s) + CO2(g)
Ca(OH)2(s)  CaO(s) + H2O(g)

62. Down the group, the size of cations increases
40.3 Variation in Properties of the compounds of the s-Block Elements (p.62)
Down the group, the size of cations increases
 polarizing power decreases
 compound with large anion become more stable
∴ thermal stability of carbonates & hydroxides of Groups I and II metals increases down the group
Effect of sizes of cations on thermal stability of compounds
Do Q. 2b on p. 73

63. Q. Explain briefly why lithium hydrogencarbonate
Q. Explain briefly why lithium hydrogencarbonate does not exist as a solid while other Group I hydrogencarbonates can be found in solid state.
A. In solid form, the cation and anion are close to each other. Due to small size of Li+, it has a high polarizing power. This distorts the electron cloud of HCO3-, making the anion unstable.
As the size of cations increases down the group, the polarizing power decreases, therefore, solid hydrogencarbonates can be formed.

64. Effect of Heat on s-block carbonates and hydroxides (p.19)
i. Carbonates
Group I: All are thermally stable except Lithium.
Group II: All decompose on heating forming metal oxides and carbon dioxide.
ii. Hydroxides (p.21)
Group I: All are thermally stable except Lithium.
Group II: All decompose on heating forming metal oxides and water.

65. Processes involved in Dissolution and their Energetics
40.3 Variation in Properties of the compounds of the s-Block Elements (p.63)
notes p. 21
Relative Solubility of the Sulphates(VI) and Hydroxides
Processes involved in Dissolution and their Energetics
When an ionic solid is dissolved in water, two processes are taken place: 1. Breakdown of the ionic solid (-ve lattice enthalpy) 2. Stabilization of ions by water molecules (hydration enthalpy released)

66. Dissolution of NaCl

67. Solubility of s-block Sulphates and Hydroxides (p.23)
DHs
MX(s)
M+(aq) + X-(aq)
-DU
DHhyd
M+(g) + X-(g)
A low modulus of lattice enthalpy and a high modulus of hydration enthalpy favour the dissolving process.

68. Effect of charge and size of ions on Hhyd and Hlattice

69. Solubility of s-block Sulphates and Hydroxides
i. For large anions, like sulphates
When moving down the group, the decrease in size of the cation does not cause a significant change of DU. However, DHhyd become less negative and has a significant change  the solubility of sulphates decreases down the group.
SO42-
MgSO4
SrSO4

70. ii. For smaller anions, like hydroxides
When moving down the group, the increase in size of the cation causes a significant change of DU but DHhyd change a little because of the great hydration energy of the anion. Therefore the solubility of hydroxide increases down the group.
Mg(OH)2
Sr(OH)2
iii. Group I sulphates and hydroxides are more soluble than that of Group II. Why?

71.  The enthalpy changes of solution are more –ve
40.3 Variation in Properties of the compounds of the s-Block Elements (p.23)
Relative Solubility of the Sulphates(VI) and Hydroxides –Trend and Interpretation
The sulphates(VI) and hydroxides of Group I metals are more soluble in water than those of Group II metals
∵ Group I metals has a smaller charge and larger size than Group II metals in the same period
 The lattice enthalpies of Group I compounds are smaller in magnitude than those of Group II compounds
 The enthalpy changes of solution are more –ve

72. Do Q. 6, 10 and Q. 7 on p. 74
The END