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The s-Block Elements
1
The s-Block Elements
•
Elements of Groups IA* (the alkali metals)
and IIA* (the alkaline earth metals)
 constitute the s-block elements
 their outermost shell electrons are in
the s orbital
*Note: In the following, Groups IA and IIA are
abbreviated as Groups I and II respectively.
2
3
The s-block elements
The s-Block Elements
•
Similarities
1. highly reactive metals
2. strong reducing agents
3. form ionic compounds with fixed
oxidation states of +1 for Group I
elements and +2 for Group II elements
4
Q.1
Group I
Li
Lithium
Na
Sodium
K
Potassium
Rb
Rubidium
Cs
Caesium
*Fr
Francium
Electronic
configuration
[He] 2s1
[Ne] 3s1
[Ar] 4s1
[Kr] 5s1
[Xe] 6s1
[Rn] 7s1
Ca
Calcium
Sr
Strontium
Ba
Barium
*Ra
Radium
[Ar] 4s2
[Kr] 5s2
[Xe] 6s2
[Rn] 7s2
Group II
Electronic
configuration
5
Be
Mg
Beryllium Magnesium
[He] 2s2
[Ne] 3s2
Group I elements
•
6
Lithium
Group I elements
•
7
Sodium
Group I elements
•
8
Potassium
Group I elements
•
9
Rubidium
Group I elements
•
10
Francium - radioactive
Group I elements
•
11
Beryllium
Group I elements
•
12
Magnesium
Group I elements
•
13
Calcium
Group I elements
•
14
Strontium
Group I elements
•
15
Barium
Group I elements
•
16
Radium - radioactive
Characteristic
Properties of the
s-Block Elements
17
Group I
element
Electronegativity
value
Group II
element
Electronegativity
value
Li
1.0
Be
1.5
Na
0.9
Mg
1.2
K
0.8
Ca
1.0
Rb
0.8
Sr
1.0
Cs
0.7
Ba
0.9
Fr
–
Ra
–
All have low electronegativity.
 electropositive
18
Group I
element
Electronegativity
value
Group II
element
Electronegativity
value
Li
1.0
Be
1.5
Na
0.9
Mg
1.2
K
0.8
Ca
1.0
Rb
0.8
Sr
1.0
Cs
0.7
Ba
0.9
Fr
–
Ra
–
EN  down the group
EN : Group II > Group I (∵ greater ENC)
19
Group I
m.p.(C)
b.p.(C) Group II m.p.(C)
b.p.(C)
Li
181
1342
Be
1287
2469
Na
98
883
Mg
650
1090
K
63
760
Ca
850
1492
Rb
39
688
Sr
770
1367
Cs
29
690
Ba
714
1637
Fr
-
-
Ra
-
-
Bonding
Strength of metallic bond : Group II > Group I
m.p./b.p. : Group II > Group I
20
Group I
m.p.(C)
b.p.(C) Group II m.p.(C)
Li
181
1342
Be
1287
2469
Na
98
883
Mg
650
1090
K
63
760
Ca
850
1492
Rb
39
688
Sr
770
1367
Cs
29
690
Ba
714
1637
Fr
-
-
Ra
-
-
Hardness : Group I < Group II
Na/K…can be easily cut with a knife
21
b.p.(C)
Density
Group I
Structure
Li
b.c.c.
0.53
Be
h.c.p.
1.86
Na
b.c.c.
0.97
Mg
h.c.p.
1.74
K
b.c.c.
0.86
Ca
f.c.c.
1.55
Rb
b.c.c.
1.53
Sr
f.c.c.
2.54
Cs
b.c.c.
1.90
Ba
b.c.c.
3.59
Fr
-
-
Ra
-
-
(g
cm3)
Group II Structure
Density
(g cm3)
Structure
Group I : b.c.c. Group II : f.c.c. or h.c.p. except Ba
Density : Group II > Group I
22
Density
Group I
Structure
Li
b.c.c.
0.53
Be
h.c.p.
1.86
Na
b.c.c.
0.97
Mg
h.c.p.
1.74
K
b.c.c.
0.86
Ca
f.c.c.
1.55
Rb
b.c.c.
1.53
Sr
f.c.c.
2.54
Cs
b.c.c.
1.90
Ba
b.c.c.
3.59
Fr
-
-
Ra
-
-
(g
cm3)
Group II Structure
Density
(g cm3)
Structure
Group I : b.c.c. Group II : f.c.c. or h.c.p. except Ba
Density also depends on size and mass of the atoms
23
Group I
E o (V)
Group II
E o (V)
Li
-3.04
Be
-1.69
Na
-2.72
Mg
-2.37
K
-2.92
Ca
-2.87
Rb
-2.99
Sr
-2.89
Cs
-3.02
Ba
-2.90
Metallic charater (Reactivity) : High tendency to lose electrons as shown by –ve E
Mn+(aq) + ne  M(s)
24
Group I
E o (V)
Group II
Li
-3.04
Be
-1.69
Na
-2.72
Mg
-2.37
K
-2.92
Ca
-2.87
Rb
-2.99
Sr
-2.89
Cs
-3.02
Ba
-2.90
Metallic charater (Reactivity) : -
 down the groups
Group I > Group II
25
E o (V)
sodium
Sodium is stored under paraffin oil
26
caesium
rubidium
Caesium and rubidium are stored
in vacuum-sealed ampoules
27
Formation of Basic Oxides
1. Group I Elements
•
28
All alkali metals form more than one
type of oxide on burning in air (except
lithium)
1. Group I Elements
•
Three types of oxides:
 normal oxides
 peroxides
 superoxides
Abundant supply
O2–
oxide
ion
29
1
O2
2
O
 O22– 
2O2–
2
peroxide
ion
superoxide
ion
1. Group I Elements
Type of oxide formed depends on
1. supply of oxygen
2. reaction temperature
30
3. charge density of M+
1. Group I Elements
•
Lithium
 when it is burnt in air, it forms normal
oxide only
 2Li2O(s)
4Li(s) + O2(g) 
180 C
lithium oxide
31
1. Group I Elements
•
Sodium
 when it is burnt in an abundant
supply of oxygen
 forms both the normal oxide and the
peroxide
 2Na2O(s)
4Na(s) + O2(g) 
180 C
sodium oxide
2Na2O(s) + O2(g)  2Na2O2(s)
300 C
excess
32
sodium peroxide
1. Group I Elements
•
Potassium, rubidium and caesium
 form All three types of oxides when
burnt in sufficient supply of oxygen
33
1. Group I Elements
•
Potassium:
4K(s) + O2(g)  2K2O(s)
potassium oxide
2K2O(s) + O2(g)  2K2O2(s)
potassium peroxide
K2O2(s) + O2(g)  2KO2(s)
potassium superoxide
34
1. Group I Elements
•
Rubidium:
4Rb(s) + O2(g)  2Rb2O(s)
2Rb2O(s) + O2(g)  2Rb2O2(s)
Rb2O2(s) + O2(g)  2RbO2(s)
35
1. Group I Elements
•
Caesium:
4Cs(s) + O2(g)  2Cs2O(s)
2Cs2O(s) + O2(g)  2Cs2O2(s)
Cs2O2(s) + O2(g)  2CsO2(s)
36
Oxides formed by Group I elements
Group I
element
Normal oxide
Peroxide
Superoxide
Li
Li2O
–
–
Na
Na2O
Na2O2
–
K
K2O
K2O2
KO2
Rb
Rb2O
Rb2O2
RbO2
Cs
Cs2O
Cs2O2
CsO2
Cations with high charge densities (Li+ or Na+) tend to
polarize the large electron clouds of peroxide ions and/or
superoxide ions
 Making them decompose to give oxide ions
37
1. Group I Elements
The electron cloud of the superoxide ion is
greatly distorted by the small lithium ion
38
Oxides formed by Group I elements
Group I
element
Normal oxide
Peroxide
Superoxide
Li
Li2O
–
–
Na
Na2O
Na2O2
–
K
K2O
K2O2
KO2
Rb
Rb2O
Rb2O2
RbO2
Cs
Cs2O
Cs2O2
CsO2
White
solids
Slightly
coloured
solids
Highly
coloured
solids
39
KO2 used as oxygen generators and CO2
scrubbers in spacecrafts
4KO2 + 2H2O  4KOH + 3O2
2KOH + CO2  K2CO3 + H2O
40
2. Group II Elements
•
Beryllium, magnesium and calcium
 form normal oxides only on burning in
air
2Be(s) + O2(g)  2BeO(s)
2Mg(s) + O2(g)  2MgO(s)
2Ca(s) + O2(g)  2CaO(s)
41
Q.2(a)
Be2+, Mg2+ and Ba2+ have higher charge
densities
 more polarizing
 distort the electron cloud of O22
 O22 decomposes to give O2
42
Q.2(b)
2Sr(s) + O2(g)  2SrO(s)
strontium oxide
2SrO(s) + O2(g)
Sr(s) + O2(g)
43
2SrO2(s)
strontium peroxide
SrO2(s)
Q.2(b)
2Ba(s) + O2(g)  2BaO(s)
barium oxide
2BaO(s) + O2(g)
500C
700C
2BaO2(s)
barium peroxide
Ba(s) + O2(g)
44
BaO2(s)
Oxides formed by Group II elements
Group II
Normal oxide
element
Peroxide
Superoxide
Be
BeO
–
–
Mg
MgO
–
–
Ca
CaO
–
–
Sr
SrO
SrO2
–
Ba
BaO
BaO2
–
KO2
superoxide
45
Oxides formed by Group II elements
Group II
Normal oxide
element
Peroxide
Superoxide
Be
BeO
–
–
Mg
MgO
–
–
Ca
CaO
–
–
Sr
SrO
SrO2
–
Ba
BaO
BaO2
–
All these oxides are basic in nature (except
beryllium oxide which is amphoteric)
46
Formation of hydroxides
1. Group I hydroxides
2Li(s) + 2H2O(l)  2LiOH(aq) + H2(g)
2Na(s) + 2H2O(l)  2NaOH(aq) + H2(g)
2K(s) + 2H2O(l)  2KOH(aq) + H2(g)
2Rb(s) + 2H2O(l)  2RbOH(aq) + H2(g)
2Cs(s) + 2H2O(l)  2CsOH(aq) + H2(g)
47
Formation of hydroxides
1. Group I hydroxides
For normal oxides,
M2O(s) + H2O(l)  2MOH(aq)
For peroxides,
M2O2(s) + 2H2O(l)  2MOH(aq) + H2O2(aq)
For superoxides,
2MO2(s) + 2H2O(l)  2MOH(aq) + H2O2(aq) + O2(g)
48
Formation of hydroxides
2. Group II hydroxides
Ca(s) + 2H2O(l)  Ca(OH)2(aq) + H2(g)
Sr(s) + 2H2O(l)  Sr(OH)2(aq) + H2(g)
Ba(s) + 2H2O(l)  Ba(OH)2(aq) + H2(g)
Mg reacts with steam but not water.
Mg(s) + H2O(g)  MgO(s) + H2(g)
Be does not react with water and steam.
49
Formation of hydroxides
2. Group II hydroxides
CaO(s) + H2O(l)  Ca(OH)2(aq)
SrO(s) + H2O(l)  Sr(OH)2(aq)
BaO(s) + H2O(l)  Ba(OH)2(aq)
MgO(s) + H2O(l)
slightly soluble
Mg(OH)2(aq)
BeO(s) + H2O(l)  No reaction
50
Ionic Bonding with Fixed
Oxidation State in their Compounds
Predominantly ionic
Group I : +1
∵ Low 1st I.E. but very high 2nd I.E.
Group II : +2
∵ Low 1st and 2nd I.E. but very high 3rd I.E.
51
Chemical formulae of some Group I compounds and the
oxidation states of Group I elements in the compounds
Group I
element
Oxide
Hydride
Chloride
Oxidation state of Group
I element in the compound
Li
Li2O
LiH
LiCl
+1
Na
Na2O2
NaH
NaCl
+1
K
KO2
KH
KCl
+1
Rb
RbO2
RbH
RbCl
+1
Cs
CsO2
CsH
CsCl
+1
52
Chemical formulae of some Group II compounds and the
oxidation states of Group II elements in the compounds
Group II
element
Oxide
Hydride
Chloride
Oxidation state of Group
II element in the
compound
Be
BeO
BeH2
BeCl2
+2
Mg
MgO
MgH2
MgCl2
+2
Ca
CaO
CaH2
CaCl2
+2
Sr
SrO
SrH2
SrCl2
+2
Ba
BaO
BaH2
BaCl2
+2
53
Weak Tendency to Form Complexes
A complex is formed when a central metal
atom or ion is surrounded by other
molecules or ions (called ligands) which
form dative covalent bonds with the central
metal atom or ion using their lone pair.
54
Weak Tendency to Form Complexes
Unlike transition metals, all s-block metals
(except Be) show little tendency to form
complexes
55
Weak Tendency to Form Complexes
Reasons : 1. Absence of low-lying vacant d-orbtals to
accept lone pairs from ligands.
For Na+, 1s2, 2s2, 2p6, 3s, 3p, 3d
High-lying relative to 2p
For Fe2+, 1s2, 2s2, 2p6, 3s2, 3p3, 3d6
Low-lying relative to 3p
56
Weak Tendency to Form Complexes
Reasons : 2. s-block cations (M+, M2+) have relatively
low charge densities
 less polarizing and less able to accept
lone pairs from ligands.
57
A complex ion, [Co(NH3)6]3+
NH 3
NH 3
H3N
Co3+
NH 3
H3N
All six bonds are strong dative
covalent bonds
NH 3
A hydrated ion, Na+(aq)
OH 2
H2O
Na
+
H2O
OH 2
58
OH 2
Dipole-ion attraction
OH 2
Weaker than dative bond
Weak Tendency to Form Complexes
Owing to its high charge density,
Be2+ can form complexes
59
[Be(H2O)4]2+(aq) + H2O(l)
[Be(H2O)3(OH)]+(aq) + H2O(l)
[Be(H2O)2 (OH)2](s) + H2O(l)
[Be(H2O)(OH)3 ] (aq) + H2O(l)
[Be(H2O)3(OH)]+(aq) + H3O+(aq)
[Be(H2O)2(OH)2](s) + H3O+(aq)
[Be(H2O)(OH)3](aq) + H3O+(aq)
[Be(OH)4]2(aq) + H3O+(aq)
Overall reaction : (1) + (2) + (3) + (4)
[Be(H2O)4]2+(aq) + 4H2O(l)
60
[Be(OH)4]2(aq) + 4H3O+(aq)
[Be(H2O)4]2+(aq) + H2O(l)
[Be(H2O)3(OH)]+(aq) + H2O(l)
[Be(H2O)2 (OH)2](s) + H2O(l)
[Be(H2O)(OH)3 ] (aq) + H2O(l)
[Be(H2O)3(OH)]+(aq) + H3O+(aq) (1)
[Be(H2O)2(OH)2](s) + H3O+(aq) (2)
[Be(H2O)(OH)3](aq) + H3O+(aq) (3)
[Be(OH)4]2(aq) + H3O+(aq) (4)
Overall reaction : (1) + (2) + (3) + (4)
[Be(H2O)4]2+(aq) + 4H2O(l)
[Be(OH)4]2(aq) + 4H3O+(aq)
pH   equilibrium positions shifts to the right
61
[Be(H2O)4]2+(aq) + H2O(l)
[Be(H2O)3(OH)]+(aq) + H2O(l)
[Be(H2O)2 (OH)2](s) + H2O(l)
[Be(H2O)(OH)3 ] (aq) + H2O(l)
[Be(H2O)3(OH)]+(aq) + H3O+(aq) (1)
[Be(H2O)2(OH)2](s) + H3O+(aq) (2)
[Be(H2O)(OH)3](aq) + H3O+(aq) (3)
[Be(OH)4]2(aq) + H3O+(aq) (4)
(1) + (2)
[Be(H2O)4]2+(aq) + 2H2O(l)
[Be(H2O)2(OH)2](s) + 2H3O+(aq)
+ 2OH(aq)
[Be(H2O)4]2+(aq) + 2OH(aq)
+ 2OH(aq)
[Be(H2O)2(OH)2](s) + 2H2O
Or simply,
Be2+(aq) + 2OH(aq)
62
Be(OH)2(s)
Characteristic Flame Colours of Salts
Most s-block elements and their compounds give
a characteristic flame colour in the flame test
Group I
element
Li
Na
K
Rb
Cs
63
Flame colour
Crimson
Golden yellow
Lilac
Bluish red
Blue
Group II
element
Be
Mg
Ca
Sr
Ba
Flame colour
Bright white
Brick red
Blood red
Apple green
Mechanism : 1. In the hotter part of the flame,
Na(g)
heat
Ground state
[Ne] 3s1
Na(g)*
[Ne] 3p1
2. In the cooler part of the flame,
Na(g)*
[Ne] 3p1
64
cool
Na(g) + golden yellow light
[Ne] 3s1
Visible
region
Mechanism : For salts of s-block elements,
the metal ions of the salts are first converted to
metal atoms
Na2CO3(s)
Na+Cl
65
heat
Na(g)
heat
Na(g)*
cool
Conc. HCl
Na+Cl (more volatile)
Na(g) + Cl(g)
Na(g)*
Na(g) + golden yellow light
Q.3
heat
Na+(g)
[He] 2s2 2p6
Na+(g)*
Na+(g)*
[He] 2s2 2p5 3s1
cool
[He] 2s2 2p5 3s1
Na+(g)
+ uv light
[He] 2s2 2p6
3p
visible
3s
uv
2p
66
Li
Na
K
Ca
Pt or nichrome(an alloy of Ni and Cr) is suitable
for making the wire because
1.They have no reaction with conc. HCl
2.They do not impart visible light when heated
67
Variation in Physical Properties of
s-block Elements
1. Atomic Radius and Ionic Radius
2. Ionization Enthalpies
3. Hydration Enthalpies
4. Melting Points
68
1. Atomic Radius and Ionic Radius
Group I
element
Atomic radius
(nm)
Group II
element
Atomic radius
(nm)
Li
0.152
Be
0.112
Na
0.186
Mg
0.160
K
0.231
Ca
0.197
Rb
0.244
Sr
0.215
Cs
0.262
Ba
0.217
Fr
0.270
Ra
0.220
 down the groups
∵ the outermost electrons are further away from
the nuclei
69
1. Atomic Radius and Ionic Radius
Group I
element
Atomic radius
(nm)
Group II
element
Atomic radius
(nm)
Li
0.152
Be
0.112
Na
0.186
Mg
0.160
K
0.231
Ca
0.197
Rb
0.244
Sr
0.215
Cs
0.262
Ba
0.217
Fr
0.270
Ra
0.220
Group II < Group I
∵ ENC  from left to right across the periods
70
71
On moving down the groups,
first  sharply (e.g. from Li to K)
then slowly (e.g. from K to Fr)
1. There is a sharp  in NC from 19K to 37Rb
Outermost e is drawn closer to the nucleus
72
73
2. The inner d-electrons (of Rb, Cs, Sr, Ba)
have poor shielding effect on the outermost
electrons  transition contraction
2. Ionization Enthalpy
Group I
element
1st
IE
2nd
IE
Group II
element
1st IE
2nd IE
3rd IE
Li
519
7 300
Be
900
1 760
14 800
Na
494
4 560
Mg
736
1 450
7 740
K
418
3 070
Ca
590
1 150
4 940
Rb
402
2 370
Sr
548
1 060
4 120
Cs
376
2 420
Ba
502
966
3 390
Fr
381
–
Ra
510
979
–
Both atomic radius and ENC  down the groups
Atomic radius is more important
IE  down the groups
74
2. Ionization Enthalpy
Group I
element
1st
IE
2nd
IE
Group II
element
1st IE
2nd IE
3rd IE
Li
519
7 300
Be
900
1 760
14 800
Na
494
4 560
Mg
736
1 450
7 740
K
418
3 070
Ca
590
1 150
4 940
Rb
402
2 370
Sr
548
1 060
4 120
Cs
376
2 420
Ba
502
966
3 390
Fr
381
–
Ra
510
979
–
For Group I elements, 2nd IE >> 1st IE because
1.the outer s-electron is well shielded by inner
shell electrons
75
2. Ionization Enthalpy
Group I
element
1st
IE
2nd
IE
Group II
element
1st IE
2nd IE
3rd IE
Li
519
7 300
Be
900
1 760
14 800
Na
494
4 560
Mg
736
1 450
7 740
K
418
3 070
Ca
590
1 150
4 940
Rb
402
2 370
Sr
548
1 060
4 120
Cs
376
2 420
Ba
502
966
3 390
Fr
381
–
Ra
510
979
–
For Group I elements, 2nd IE >> 1st IE because
2. the 2nd electron is closer to the nucleus and
is poorly shielded by other electrons in the
76
same shell which is completely filled.
2. Ionization Enthalpy
Group I
element
1st
IE
2nd
IE
Group II
element
1st IE
2nd IE
3rd IE
Li
519
7 300
Be
900
1 760
14 800
Na
494
4 560
Mg
736
1 450
7 740
K
418
3 070
Ca
590
1 150
4 940
Rb
402
2 370
Sr
548
1 060
4 120
Cs
376
2 420
Ba
502
966
3 390
Fr
381
–
Ra
510
979
–
For Group II elements, 3rd IE >> 2nd IE
Similar reasons can be applied
77
Variations in the first and second ionization
enthalpies of Group I elements
78
Variations in the first,
second and third
ionization enthalpies of
Group II elements
79
2. Ionization Enthalpy
Group I
element
1st
IE
2nd
IE
Group II
element
1st IE
2nd IE
3rd IE
Li
519
7 300
Be
900
1 760
14 800
Na
494
4 560
Mg
736
1 450
7 740
K
418
3 070
Ca
590
1 150
4 940
Rb
402
2 370
Sr
548
1 060
4 120
Cs
376
2 420
Ba
502
966
3 390
Fr
381
–
Ra
510
979
–
Group II > Group I ∵
The outer s-electrons of Group II atoms are
closer to the nucleus and experience higher ENC
80
3. Hydration enthalpy
Hydration enthalpy (Hhyd) is the amount of
energy released when one mole of aqueous ions
is formed from its gaseous ions.
M+(g) + aq  M+(aq)
H = Hhyd
M2+(g) + aq  M2+(aq)
H = Hhyd
 always has a negative value
81
Group I ion
Hydration enthalpy
(kJ mol–1)
Group II ion
Hydration
enthalpy (kJ mol–1)
Li+
–519
Be 2+
–2 450
Na+
–406
Mg2+
–1 920
K+
–322
Ca2+
–1 650
Rb+
–301
Sr2+
–1 480
Cs+
–276
Ba2+
–1 360
Fr+
–
Ra2+
–
 down the groups
∵ charge density of metal ions  down the groups
 attraction between ions and water molecules 
+
82
Group I ion
Hydration enthalpy
(kJ mol–1)
Group II ion
Hydration
enthalpy (kJ mol–1)
Li+
–519
Be 2+
–2 450
Na+
–406
Mg2+
–1 920
K+
–322
Ca2+
–1 650
Rb+
–301
Sr2+
–1 480
Cs+
–276
Ba2+
–1 360
Fr+
–
Ra2+
–
Group II > Group I
∵ Group II ions have higher charge and small size
 higher charge density
 stronger ion-dipole interaction
83
Variations in hydration
enthalpy of the ions of
Groups I and II
elements
84
4. Melting Point
The melting points of s-block elements depend
on the metallic bond strength which in turn
depends on
1. charge density of cations
2. number of valence electrons participating
in the sea of electrons
3. packing efficiency of the crystal lattices
85
Group I
element
Melting Point
(C)
Group II
element
Melting Point
(C)
Li
180
Be
1280
Na
97.8
Mg
650
K
63.7
Ca
850
Rb
38.9
Sr
768
Cs
28.7
Ba
714
Fr
24
Ra
697
1.  down the groups
∵ ionic radii  down the groups
 charge density 
 interaction between ions and electron sea 
86
Group I
element
Melting Point
(C)
Group II
element
Melting Point
(C)
Li
180
Be
1280
Na
97.8
Mg
650
K
63.7
Ca
850
Rb
38.9
Sr
768
Cs
28.7
Ba
714
Fr
24
Ra
697
2. Group II > Group I ∵
(a)Group II cations have higher charge density
(b)
More valence electrons are involved in
the sea
of electrons
(c)Packing efficiency : Group II > Group I
87
Reason not known !!
88
Variation in Chemical Properties
s-Block elements have strong reducing power
∵
low ionization enthalpies
low atomization enthalpies
89
M+(g)
~Ea
Ionization
enthalpy
M(g)
M(s)
Atomization
enthalpy
Hydration
enthalpy
ΔH0 reaction  0
M+(aq)
M(s)  M+(aq) + e
90
H < 0
M+(g)
~Ea
Ionization
enthalpy
M(g)
M(s)
Atomization
enthalpy
Hydration
enthalpy
ΔH0 reaction  0
M+(aq)
Reactivity : Na > Ca (depends on Ea)
91
Position in e.c.s. : Ca > Na (depends on Ho or Eo
Variation in Chemical Properties
The reactivity of s-block elements  down the
groups
∵ both I.E. and A.E.  down the groups
 Ea  down the groups
 Reaction rate  down the groups
92
Variation in Chemical Properties
Reactivity : Group I > Group II
∵ both I.E. and A.E.  across the periods
 Ea  across the periods
 Reaction rate  across the periods
93
1. Reactions with hydrogen
Group I
2M(s) + H2(g)
Group II
M(s) + H2(g)
94
300C – 500C
600C – 700C
2MH(s)
MH2(s)
1. Reactions with hydrogen
4LiH + AlCl3
Dry ether
LiAlH4 + 3LiCl
Reducing agent in
organic syntheses
95
2. Reactions with Oxygen
Most s-block elements

show a silvery white lustre when they
are freshly cut

they tarnish rapidly upon exposure to
the atmosphere

96
they react with oxygen in the air
to form an oxide layer
Sodium shows a silvery white lustre when freshly cut
97
Group I (p.2)
Group II
2M(s) + O2(g)
M(s) + O2(g)
98
heat
heat
2MO(s)
MO2(s)
3. Reactions with Chlorine
Group I
2M(s) + Cl2(g)
Group II
M(s) + Cl2(g)
99
heat
heat
2MCl(s)
MCl2(s)
4. Reactions with water or steam
Group I
2M(s) + H2O(l)
Group II
M(s) + 2H2O(l)
heat
heat
2MOH(aq) + H2(g)
M(OH)2(aq) + H2(g)
Mg reacts with steam but not water
Mg(s) + H2O(g)
100
heat
MgO(s) + H2(g)
Be has no reaction with either water or steam
Variation in chemical properties of the
compounds of s-block elements
Reactions of oxides
Reactions of hydrides
Reactions of chlorides
101
Reactions of oxides
1. Reactions with water
Group I
M2O(s) + H2O(l)  2MOH(aq)
M2O2(s) + 2H2O(l)  2MOH(aq) + H2O2(aq)
2MO2(s) + 2H2O(l)  2MOH(aq) + H2O2(aq) + O2(g)
102
Na2O2 is used in qualitative analysis of Cr3+
green
2Cr(OH)3(s) + 3Na2O2(s)
 2Na2CrO4(aq) + 2NaOH(aq) + 2H2O(l)
yellow
103
Reactions of oxides
Group II
CaO(s) + H2O(l)  Ca(OH)2(aq)
SrO(s) + H2O(l)  Sr(OH)2(aq)
BaO(s) + H2O(l)  Ba(OH)2(aq)
MgO(s) + H2O(l)
slightly soluble
Mg(OH)2(aq)
BeO(s) + H2O(l)  No reaction
104
increasing
basicity
Reactions of oxides
2. Reactions with acids
Group I
M2O(s) + 2HCl(aq)  2MCl(aq) + H2O(l)
M2O2(s) + 2HCl(aq)  2MCl(aq) + H2O2(aq)
2MO2(s) + 2HCl(aq)  2MCl(aq) + H2O2(aq) + O2(g)
Group II More vigorous than those with water
MO(s) + 2HCl(aq)  MCl2(aq) + H2O(l)
105
Reactions of oxides
3. Reactions with alkalis
Reaction with water instead except BeO
BeO(s) + 2OH(aq) + H2O(l)  Be(OH)42(aq)
amphoteric
106
Reactions of hydrides
MOH(aq) + H2(g)
MH(s)
MCl(aq) + H2(g)
H (a strong base) tends to react with protonic
reagents to release H2
Reactivity  down the groups
107
Reactions of chlorides
Group I
No significant reactions with water, acids or alkalis
Group II
Do not undergo significant hydrolysis except
BeCl2 and MgCl2 More favoured in alkaline solutions
BeCl2(aq) + 2H2O(l)  Be(OH)2(aq) + 2HCl(aq)
MgCl2(aq) + H2O(l)  Mg(OH)Cl(aq) + HCl(aq)
Basic salt
108
Relative Thermal Stability of the
Carbonates and Hydroxides of
s-Block Elements
Thermal stability refers to the resistance
of a compound to undergo decomposition
on heating.
109
Thermal decomposition reactions
Metal carbonates
M2CO3(s)
heat
M2O(s) + CO2
MCO3(s)
heat
MO(s) + CO2
Metal hydroxides
110
2MOH(s)
heat
M2O(s) + H2O(g)
M(OH)2(s)
heat
MO(s) + H2O
Relative thermal stability can be measured
in two ways
1. By comparing the decomposition
temperatures
A higher decomposition temperature
 a greater thermal stability
111
Metal
carbonate
Decomposition
temperature
/C
BeCO3
MgCO3
CaCO3
SrCO3
BaCO3
~100
540
900
1290
1360
Decomposition temperature is
the temperature at which the pressure of CO2 in
equilibrium with the solid carbonate reaches 1 atm
in a closed system.
Below the DT, some CO2 can still be detected but
the pressure is less than 1 atm
112
1. The Carbonates
•
Example:
100 C
BeCO3(s)  BeO(s) + CO2(g)
540 C
MgCO3(s)  MgO(s) + CO2(g)
900 C
CaCO3(s)  CaO(s) + CO2(g)
1290  C
SrCO3(s)   SrO(s) + CO2(g)
1360  C
BaCO3(s)   BaO(s) + CO2(g)
113
Relative thermal stability can be measured
in two ways
2. By comparing the standard enthalpy
changes of thermal decomposition
reactions
M(OH)2(s)  MO(s) + H2O(g) H > 0
A more positive H value
 a thermally more stable compound
114
Metal
hydroxide
Ho /
kJ mol1
Be(OH)2 Mg(OH)2 Ca(OH)2 Sr(OH)2 Ba(OH)2
+54
+81
+109
+127
+146
Trends : 1.
 down the groups
2.
Group I > Group II
3.
Li resembles Mg more than the other
group 1 elements (diagonal relationship,
pp.14-15)
115
2. The Hydroxides

Be(OH)2(s) 
BeO(s) + H2O(g)

H = +54 kJ mol–1

Mg(OH)2(s) 
MgO(s) + H2O(g)

H = +81 kJ mol–1

Ca(OH)2(s) 
CaO(s) + H2O(g)

H = +109 kJ mol–1

Sr(OH)2(s) 
SrO(s) + H2O(g)

H = +127 kJ mol–1
116

Ba(OH)2(s) 
BaO(s) + H2O(g)

H = +146 kJ mol–1
Factors affecting thermal stability of
carbonates and hydroxides
1.
Polarizing power of cation
2. Polarizability of polyatomic anion
3. Lattice enthalpy of metal oxide produced
117
Interpretation of trends in thermal
stability of carbonates and hydroxides
1.
Group I > Group II
(a)
M2+ ions have higher charge
densities than M+ ions
 M2+ ions are more polarizing
than M+ ions
 Can polarize more the electron
cloud of polyatomic anions
118
polarization
M2+
O
O
heat
C
MO + CO2
O
polarization
M2+
119
heat
O
H
O
H
MO + H2O
Polarizability  as the size of anion 
120
Thermal
decomposition
Polyatomic
ion
121
When a compound with large anions undergoes thermal
decomposition, a compound with small anions will be
formed since small anions are less easily polarized
122
more stable
compound
with stronger
bond
Simple ion
123
M2+
S2
Simple ion
124
polarization
M
S
Stronger ionic
bond with covalent
character
Interpretation of trends in thermal
stability of carbonates and hydroxides
1.
Group I > Group II
(b)
M2+ ions have higher charge
densities than M+ ions
 Lattice enthalpy : MO > M2O
 Energetic stability : MO > M2O
125
CaCO3(s)
Na2CO3(s)
more favourable
heat
less favourable
heat
CaO(s) + CO2(g)
more stable
Na2O(s) + CO2(g)
less stable
Thermal stability of carbonates : Group I > Group II
126
Interpretation of trends in thermal
stability of carbonates and hydroxides
2. Thermal stability  down the groups
∵ size of cations  down the groups
∴ (a) charge density/polarizing power
of cation  down the groups
(b) lattice enthalpies of MO/M2O 
down the groups
127
MgCO3(s)
more favourable
more polarized
BaCO3(s)
less polarized
heat
less favourable
heat
MgO(s) + CO2(g)
more stable
BaO(s) + CO2(g)
less stable
Thermal stability of carbonates :  down the groups
128
Effect of sizes of the cations on thermal
stability of the carbonates and hydroxides
of both Groups I and II metals
129
Interpretation of trends in thermal
stability of carbonates and hydroxides
3. Li compounds resemble Mg compounds
(diagonal relationship)
Charge density/polarizing power : Li+  Mg2+
130
Interpretation of trends in thermal
stability of carbonates and hydroxides
4. Thermal stability of nitrates follows
similar patterns (Optional)
2MNO3(s)
2M(NO3)2(s)
131
heat
heat
2MNO2 + O2
2MO + 4NO2 + O2
Relative Solubility of the
Sulphates(VI) and Hydroxides of
s-Block Elements
In general,
Group I >> Group II
132
Q.4
Compounds
Mg(OH)2
Ca(OH)2
Sr(OH)2
Ba(OH)2
Solubility / mol per
100 of water
0.02  103
1.5  103
3.4  103
15  103
Solubility / mol per
Compounds
100 of water
MgSO4
1800  104
CaSO4
11  104
SrSO4
0.71  104
BaSO4
0.009  104
133
In general,
Size and/or charge of
the anion 
 Polarizability of anion 
 Covalent character 
 Solubility in water 
Solubility / mol per
Compounds
100 of water
Mg(OH)2
0.02  103
Ca(OH)2
1.5  103
Sr(OH)2
3.4  103
Ba(OH)2
15  103
 down the group
Solubility / mol per
100 of water
1800  104
11  104
0.71  104
0.009  104
 down the group
Compounds
MgSO4
CaSO4
SrSO4
BaSO4
134
1. Processes involved in Dissolution and
their Energetics
•
Two processes are
1.
the breakdown of the ionic lattice
2. the subsequent stabilization of the
ions by water molecules (this process
is called hydration)
135
1. the breakdown of the ionic lattice
NaCl(s)  Na+(g) + Cl(g)
H1 = (lattice enthalpy) > 0
2. the subsequent stabilization of the
ions by water molecules (this process
is called hydration)
Na+(g) + Cl(g) + aq  Na+(aq) + Cl(aq)
H2 = (hydration enthalpy) < 0
136
Hsolution
NaCl(s)
Na +(aq)
+
Cl-(aq)
Na+(g) + Cl -(g)
ΔH
o
solution
 ΔH
o
hydration
 ΔH
o
lattice
= (-772 +776) kJ mol1
137
= +4 kJ mol1
o
ΔH
If
solution  0 , we expect the solids to dissolve
in water
o
Solubility  as ΔHsolution
becomes more –ve (less +ve)
o
Solids (e.g. NaCl) with small +ve ΔHsolution
values
are also soluble in water if the dissolution involves
an increase in the entropy of the system.
ΔG
o
solution
138
 ΔH
o
solution
 TS
o
solution
ΔG
o
solution
ΔG
o
solution
 ΔH
o
solution
0
 TS
o
solution
 Spontaneous dissolution
TS
o
is always positive
solution
Dissolution with slightly positive
can be spontaneous
139
H
o
solution
Trends and Interpretations
1. The solubility of Group(II) sulphate decreases down
the group
On moving down the group, cationic radius(r+) 
both
H
However,
140
o
L and
H
H
o
hydration
become less -ve
o
L  less rapidly than
H
o
hydration
Trends and Interpretations
 rSO 2  r
4
ΔH 
o
L
1
rSO 2  r
 constant
4
ΔH
o
solution
less –ve down
the group
141
 ΔH
o
hydration
less –ve down
the group
 ΔH
o
lattice
 +ve constant
 Solubility  down the group
Trends and Interpretations
 rSO 2  r
4
ΔH 
o
L
1
rSO 2  r
4
ΔH
o
solution
less –ve down
the group
142
 constant
(-ve)
 ΔH
o
hydration
 more rapidly
down the group
(+ve)
 ΔH
o
lattice
 less rapidly
down the group
 Solubility  down the group
Trends and Interpretations
2. The solubility of Group(II) hydroxides increases down
the group
On moving down the group, cationic radius(r+) 
both
H
However,
143
o
L and
H
H
o
hydration
become less -ve
o
o
Hhydration
L  more rapidly than
Trends and Interpretations
ΔH
o
solution
(-ve)
 ΔH
less +ve
more
–vedown
down
the group
o
hydration
 less rapidly
down the group
(+ve)
 ΔH
o
lattice
 more rapidly
down the group
 Solubility  down the group
144
For s-block compounds with small anions (e.g. OH, F),
solubility in water  down the group
For s-block compounds with large anions (e.g. SO42, CO32-),
solubility in water  down the group
For s-block compounds with medium size anions (e.g. Br),
solubility in water exhibits irregular pattern down the group
145
Solubility / mol per
Compounds
100 of water
5.5  101
MgBr2
CaBr2
6.3  101
SrBr2
4.3  101
BaBr2
3.3  101
Irregular
Solublily : First  and then 
ΔH
o
solution
(-ve)
 ΔH
o
hydration
Then less -ve
First
+ve Then  more
down the group rapidly
146
(+ve)
 ΔH
o
lattice
First  more
rapidly
Group II compounds with doubly-charged
anions (MX) are less soluble than those with
singly-charged anions (MY2)
Reasons :
1. HL of MX > HL of MY2
2. HL is the major factor affecting solubility
 Hsolution of MX is more positive
 Solubility : MX < MY2
147
Solubility : Group I > Group II
Reasons :
For a given anions, both HL and Hhydration become
more –ve from Group I to Group II
However, HL is the major factor affecting
solubility
 Hsolution : Group I is less positve than Group II
 Solubility : Group I > Group II
148
Diagonal relationship
149
Reaction
Other Group I
elements
Lithium
Magnesium
Combination with O2
Peroxides and
superoxides
Li2O (normal oxide)
MgO (normal oxide)
Combination with N2
No reaction
Li3N
Mg3N2
Action of heat on
carbonate
No reaction
(thermally stable)
Decomposes to give
Li2O and CO2
Decomposes to give
MgO and CO2
Action of heat on
hydroxide
No reaction
(thermally stable)
Decomposes to give
Li2O and H2O
Decomposes to give
MgO and H2O
Action of heat on
nitrate
Decomposes to give
MNO2 and O2
Decomposes to give
Li2O, NO2 and O2
Decomposes to give
MgO, NO2 and O2
Hydrogen carbonates
Exist as solids
Only exist in solution
Solubility of salts in
water
Most salts are more
soluble than those of
Li, Mg.
Fluoride, hydroxide, carbonate, phosphate,
ethanedioate are sparingly soluble.
Solubility of salts in
organic solvents.
Halides only slightly
soluble in organic
solvents
Halides (with covalent character) dissolve in
organic solvents
150
The END
151
40.1 Characteristic Properties of the s-Block Elements (SB p.40)
Metals are sometimes referred to as electropositive
elements. Why?
Answer
They have low electronegativity values.
Back
152
40.1 Characteristic Properties of the s-Block Elements (SB p.46)
s-Block compounds give a characteristic flame colour in
the flame test. Based on this, can you give one use of
s-block compounds?
Answer
s-Block compounds can be used in fireworks.
Back
153
40.1 Characteristic Properties of the s-Block Elements (SB p.48)
(a) Which ion has a greater ionic radius, potassium ion or
calcium ion? Explain your answer.
Answer
(a) Potassium ion (0.133 nm) has a greater ionic radius than calcium
ion (0.099 nm) . In fact, potassium ion and calcium ion are
isoelectronic and have the same number of electron shells.
However, calcium ion has one more proton than potassium ion,
the electron cloud of calcium ion will experience greater attractive
forces from the nucleus. This leads to a smaller ionic radius of
calcium ion.
154
40.1 Characteristic Properties of the s-Block Elements (SB p.48)
(b) Explain why Group I elements show a fixed oxidation
state of +1 in their compounds in terms of ionization
enthalpies.
Answer
155
40.1 Characteristic Properties of the s-Block Elements (SB p.48)
(b) Group I elements form ions with an oxidation state of +1 only. It is
because they have only one outermost shell electron. Once this
outermost shell electron is removed, a stable fully-filled electronic
configuration is obtained. Therefore, the first ionization enthalpies
of Group I elements are low. The second ionization involves the
removal of an electron from an inner electron shell. Once this
electron is removed, the stable electronic configuration will be
disrupted. Therefore, their second ionization enthalpies are very
high. As a result, Group I elements form predominantly ionic
compounds with non-metals by losing their single outermost shell
electron, and they form ions having a fixed oxidation state of +1.
156
40.1 Characteristic Properties of the s-Block Elements (SB p.48)
(c) Ions of Group I and Group II elements have a very low
tendency to form complexes. Give one reason to
explain your answer.
Answer
(c) As ions of Group I and Group II elements do not have low-lying
vacant orbitals available for forming dative covalent bonds with the
lone pair electrons of surrounding ligands, they rarely form
complexes.
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40.1 Characteristic Properties of the s-Block Elements (SB p.48)
(d) Give one test which would enable you to distinguish a
sodium compound from a potassium compound.
Answer
(d) Sodium compounds and potassium compounds can be
distinguished by conducting a flame test. In the flame test, sodium
compounds give a golden yellow flame, while potassium
compounds give a lilac flame.
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40.1 Characteristic Properties of the s-Block Elements (SB p.48)
What is a dative covalent bond? How is it formed?
Answer
A dative covalent bond is a covalent bond in which the shared
pair of
electrons is supplied by only one of the bonded atoms. A dative
covalent bond is formed by the overlapping of an empty orbital of an
atom with an orbital occupied by a lone pair of electrons of another
atom.
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40.2 Variation in Properties of the s-Block Elements (SB p.56)
(a) (i) List the factors that affect the value of the
ionization enthalpy of an atom.
Answer
(a) (i)
160
There are four main factors affecting the magnitude of the
ionization enthalpy of an atom. They are the electronic
configuration of an atom, the nuclear charge, the screening
effect, and the atomic radius.
40.2 Variation in Properties of the s-Block Elements (SB p.56)
(a) (ii) Why is ionization enthalpy of an atom always
positive?
Answer
(a) (ii)
161
Ionization enthalpy of an atom always has a positive value
because energy is required to overcome the attractive forces
between the nucleus and the electron to be removed.
40.2 Variation in Properties of the s-Block Elements (SB p.56)
(a) (iii) Describe the general trend of the first and second
ionization enthalpies down Group I of the
Periodic Table.
Answer
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40.2 Variation in Properties of the s-Block Elements (SB p.56)
(a) (iii) The first ionization enthalpies of Group I elements are
relatively low. The outermost s electron is located in a new
electron shell. The attractive force between this s electron and
the nucleus is relatively weak. Also, this s electron is
effectively shielded from the attraction of the nucleus by the
fully-filled inner electron shells. Once this electron is removed,
a stable octet or duplet electronic configuration is obtained.
Consequently, this s electron is relatively easy to be removed,
and hence the first ionization enthalpies of Group I elements
are relatively low. However, the second ionization of Group I
elements involves the loss of an inner shell electron which is
closer to the nucleus. The removal of this electron disrupts the
stable electronic configuration. Therefore, the second
ionization enthalpies of Group I elements are extremely high.
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40.2 Variation in Properties of the s-Block Elements (SB p.56)
(b) (i) List the factors that affect the value of the
hydration enthalpy of an ion.
Answer
(b) (i)
164
The value of the hydration enthalpy of an ion depends on the
size and the charge of the ion.
40.2 Variation in Properties of the s-Block Elements (SB p.56)
(b) (ii) Why does hydration enthalpy of an ion always
have a negative value?
Answer
(b) (ii)
165
Hydration enthalpy of an ion always has a negative value
because it is the amount of energy released resulting from the
attraction between the ion and water molecules.
40.2 Variation in Properties of the s-Block Elements (SB p.56)
(b) (iii) Describe the general trend of the hydration
enthalpy down Group II of the Periodic Table.
Answer
(b) (iii) Going down Group II, the hydration enthalpy of the
ions
decreases (becomes less negative). Since the ions get larger
in size on moving down the group, the charge density of the
ions falls. As a result, the electrostatic attraction between the
ions and water molecules becomes weaker, and the hydration
enthalpy becomes less negative down the group.
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40.2 Variation in Properties of the s-Block Elements (SB p.57)
The burning of lithium, sodium and potassium in oxygen
gives different types of oxides. Why do the metals
behave differently?
Answer
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40.2 Variation in Properties of the s-Block Elements (SB p.57)
On burning in air, lithium forms only lithium oxide, and it does not form the
peroxide or superoxide. This is because the size of lithium ion is very small,
leading to its high polarizing power. When a peroxide ion or superoxide ion
approaches a lithium ion, the electron cloud of the peroxide ion or
superoxide ion (large in size) would be greatly distorted by the lithium ion.
The greater the distortion of the electron cloud, the lower the stability of the
compound. That is why lithium peroxide and lithium superoxide do not exist.
Sodium ion has a larger size than lithium ion. Its lower polarizing power
allows it to form the peroxide when sodium is burnt in air. Potassium ion
has a much larger size, so it has relatively low polarizing power. The
electron cloud of the peroxide ion or superoxide ion would not be seriously
distorted by potassium ion. This allows the peroxide ions or superoxide ions
to pack around potassium ion with a higher stability. As a result, potassium
is able to form stable peroxide or superoxide on burning in air.
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40.2 Variation in Properties of the s-Block Elements (SB p.58)
(a) Suggest a reason why the reaction of lithium with
water is less vigorous than those of sodium and
potassium.
Answer
(a) The reactivity of Group I metals with water is related to the relative
ease of the metal atoms to lose the outermost shell electron. Going
down the group, as the atomic size increases, the outermost shell
electron becomes easier to be removed. Therefore, the reactivity of
Group I metals towards water increases down the group. Lithium
reacts with water vigorously. Sodium reacts with water violently and
moves on the water surface with a hissing sound.
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40.2 Variation in Properties of the s-Block Elements (SB p.58)
(b) Which element is the strongest reducing agent, calcium,
strontium or barium?
Answer
(b) Barium is the strongest reducing agent. It is because the reducing
power of an element is related to the ease of the atom to lose the
outermost shell electron. Since barium has larger atomic sizes, its
outermost shell electrons are less firmly held by the nucleus.
Therefore, barium has a higher tendency to lose its outermost shell
electrons than both calcium and strontium.
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40.3 Variation in Properties of the Compounds of the s-Block Elements
(SB p.64)
The value of Hsoln of a solid does not indicate whether
the solid is soluble in water or not. So how can we
predict the solubility of a solid in water?
Answer
Generally speaking, for a solid to be soluble in water, its enthalpy
change of solution has to be a negative or a small positive value.
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40.3 Variation in Properties of the Compounds of the s-Block Elements
(SB p.65)
(a) Give balanced chemical equations for the following
reactions:
(i)
Thermal decomposition of barium carbonate
(ii) Reaction between sodium peroxide and water
(iii) Reaction between calcium oxide and dilute
hydrochloric acid
Answer
(a) (i) BaCO3(s)  BaO(s) + CO2(g)

(ii) Na2O2(s) + 2H2O(l)  2NaOH(aq) + H2O2(aq)
(iii) CaO(s) + 2HCl(aq)  CaCl2(aq) + H2O(l)
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40.3 Variation in Properties of the Compounds of the s-Block Elements
(SB p.65)
(b) Suggest a reason why barium sulphate(VI) is insoluble
in water, while potassium sulphate(VI) is soluble in
water although they have cations of similar sizes and
the same anion.
(The ionic radii of potassium ion and barium ion are
0.133 nm and 0.135 nm respectively.)
Answer
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40.3 Variation in Properties of the Compounds of the s-Block Elements
(SB p.65)
(b) When an ionic solid dissolves in water, two processes are taking
place. They are the breakdown of the ionic lattice and the
subsequent stabilization of the ions by water molecules. The
enthalpy change involved in the whole dissolution process is known
as the enthalpy change of solution, Hsoln, which is equal to Hsoln =
Hhyd – Hlattice. For an ionic compound to be soluble in water, the
enthalpy change of solution has to be a negative or a small positive
value. The reason why barium sulphate(VI) is insoluble in water
while potassium sulphate(VI) is soluble in water is that potassium
ion has a smaller charge than barium ion. The Hlattice of potassium
sulphate(VI) is smaller in magnitude (less negative) than that of
barium sulphate(VI). As a result, the enthalpy change of solution of
potassium sulphate(VI) is more negative, and hence it is soluble in
water while barium sulphate(VI) is not.
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40.3 Variation in Properties of the Compounds of the s-Block Elements
(SB p.65)
(c) Compare the solubility of calcium sulphate(VI) and
barium sulphate(VI) in water. Explain your answer.
Answer
(c) Calcium sulphate(VI) is expected to be more soluble than barium
sulphate(VI). It is because calcium ion has a smaller size than
barium ion. This causes the Hhyd of calcium sulphate(VI) to be
more negative than that of barium sulphate(VI). As a result, the
Hsoln of calcium sulphate(VI) becomes more negative than that of
barium sulphate(VI), and hence calcium sulphate(VI) is more soluble
in water than barium sulphate(VI).
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