S Block Elements Jee 2i504d
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s-Block Elements
The s-block elements
s-Block Elements • Similarities • Highly reactive metals • Strong reducing agents • Form ionic compounds • Fixed oxidation state • Group I : +1 • Group II : +2
Variation in Physical Properties of s-block Elements 1. Atomic Radius and Ionic Radius 2. Ionization Enthalpies 3. Hydration Enthalpies 4. Melting Points 5. Electronegativity
4
Atomic and Ionic Radii • The atoms and ions of alkali metals are largest in their corresponding periods. • Atomic size
Li < Na < K < Rb < Cs
• Ionic Radius
Li+ < Na + < K + < Rb + < Cs +
• Atomic volume
Li < Na < K < Rb < Cs
• Charge density
Li > Na > K > Rb > Cs
5
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
down the groups ∵ the outermost electrons are further away from the nuclei Group II < Group I ∵ ENC from left to right across the periods
6
On moving down the groups, first sharply (e.g. from Li to K)
then slowly (e.g. from K to Fr) There is a sharp in NC from 19K to 37Rb Outermost e is drawn closer to the nucleus
The inner d-electrons (of Rb, Cs, Sr, Ba) have poor shielding effect on the outermost electrons transition contraction 7
Ionisation Enthalpy
8
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
Both atomic radius and ENC down the groups Atomic radius is more important IE down the groups
9
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
For Group I elements, 2nd IE >> 1st IE because the 2nd electron is closer to the nucleus and is poorly shielded by other electrons in the same shell which is completely filled.
10
For Group II elements, 3rd IE >> 2nd IE Similar reasons can be applied
Ionization Enthalpy
Variations in the first and second ionization enthalpies of Group I elements
11
Ionization Enthalpy
Variations in the first, second and third ionization enthalpies of Group II elements
12
Electronegativity
13
Electronegativity • Relatively LOW Electronegativity • These metals have more tendency to lose electron rather than to gain an electron.
• The electronegativity values decreases down the group from Li to Cs
Li > Na > K > Rb > Cs
14
ELECTRONEGATIVITY 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
All have low electronegativity => Electropositive EN down the group EN : Group II > Group I (∵ greater ENC)
Hydration
16
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
17
always has a negative value
Hydration energies • Group I • Alkali metal ions are highly hydrated. • The smaller the ionic size, the higher the degree of hydration. • Primary and secondary shell of hydration • Li ion is very small, it is heavily hydrated. • Li ion is tetrahedrally surrounded by four water molecules using its four sp3
hybrid • Group 2
• They have higher hydration energies than Alkali metals due to smaller sizes 18
Hydration energies • In aqueous solutions, degree of hydration decreases from Li+ to Cs+ due to increase in size • Ionic radii of hydrated alkali metal ions also decreases from Li+ to Cs+
• Formation of hydrated salts : − Li > Na > K Salts. − Rb and Cs salts are not hydrated • Ionic Mobility : Cs+ > Rb+ > K + > Na + > Li +
19
Hydration energies Hydration Group I ion
Enthalpy (kJ mol–1)
Hydration Group II ion
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
20
Group II > Group I ∵ Group II ions have higher charge and small size higher charge density stronger ion-dipole interaction
Melting & Boiling 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
21
Group I
Melting
Group II
Melting
element
Point (C)
element
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
down the groups ∵ ionic radii down the groups charge density interaction between ions and electron sea 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 22
STRUCTURE Group I
Densities of Li, Na, K are lesser than that of water
Densit y (g cm3)
Group II
Densit y (g cm3)
Li
0.53
Be
1.86
Na
0.97
Mg
1.74
K
0.86
Ca
1.55
Rb
1.53
Sr
2.54
Cs
1.90
Ba
3.59
Fr
-
Ra
-
Density • Alkali metals have low density. − The reason this is that they have large atomic sizes.
• Density gradually increases on moving down the group from Li to Cs Li < Na < K < Rb < Cs.
• Anomaly: K is lighter than Na
24
Effect of light • Alkali metals when irradiated with light emit electrons with ease due to low ionization enthalpies. • This phenomenon is used in photoelectric cells, particularly caesium and potassium are used as electrodes in photoelectric cells.
25
Flame Colouration
26
Flame Colouration Most s-block elements and their compounds give a characteristic flame colour in the flame test Group I element Li Na K Rb Cs
27
Flame colour Crimson Golden yellow Lilac Bluish red Blue
Flame Colouration Most s-block elements and their compounds give a characteristic flame colour in the flame test Group II element Be Mg Ca Sr Ba
28
Flame colour Brick red Blood red Apple green
Ca, Sr, Ba
Beryllium and magnesium atoms are smaller and their electrons being strongly bound to the nucleus are not excited to higher-energy levels.
Flame Colouration 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)* 29
[Ne] 3p1
cool
Na(g) [Ne] 3s1
+ golden yellow light Visible region
Flame Colouration Mechanism : For salts of s-block elements, the metal ions of the salts are first converted to metal atoms Na2CO3(s) Na+Cl
Na(g) 30
Na(g)*
Conc. HCl
heat heat cool
Na+Cl (more volatile)
Na(g) + Cl(g)
Na(g)* Na(g)
+ golden yellow light
Complex Formation
31
Complex formation • In order to form complex compounds, a metal must possess the following characteristics. − Small size − High effective nuclear charge − Tendency to accept electrons (i.e., presence of vacant orbitals) • Since alkali metals have none of these characteristics they have little tendency to form complexes. • Lithium and Beryllium forms certain complexes. (Due to their small sizes) • The complex forming tendency fall markedly down the groups as the atomic size increases
32
Complex formation : Weak Tendency 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
33
Complex formation : Weak Tendency Reasons
2. s-block cations (M+, M2+) have relatively low charge densities less polarizing and less able to accept lone pairs from ligands.
34
Complex formation : Owing to its high charge density, Be2+ can form complexes
35
Electropositive Character • The electropositive character increases down the group from Li to Cs because ionization enthalpy
decreases down the group Li > Na > K > Rb > Cs.
Eo 36
Group I
(V)
Group II
(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) down the groups Group I > Group II
Reducing Property • Powerful reducing agents • Li > Na < K = Rb > Cs (E0) • Reasons − Heat of sublimation − Ionisation enthalpy − Hydration energy
37
Reaction with Hydrogen
38
Reaction with Hydrogen Alkali metals react with hydrogen to form ionic hydrides M+H-. The reaction of alkali metals with hydrogen decreases from Li to Cs
Group I 2M(s) + H2(g)
300C – 500C
2MH(s)
Group 1 : Hydrides • The order and reactivity with hydrogen • Li > Na > K > Rb > Cs • The ionic character of the bonds in these hydrides Increases from LiH to CsH • LiH < NaH < KH < RbH < CsH
• Stability • LiH > NaH > KH > RbH > CsH
40
LAH
Dry ether
4LiH + AlCl3
LiAlH4 + 3LiCl
• Powerful reducing agent • Tetrahedral • Selective reducing agent − Reduces carbonyl compounds to alcohols.
• It reacts violently with water, so it is necessary to use absolutely dry organic solvents • Also reduces several inorganic substances
41
Sodium tetrahydridoborate (sodium borohydride)
• NaBH4 • Can be used even in aqueous solutions • Na and K hydrides are useful
42
Reaction with Hydrogen
Alkaline earth metals react with hydrogen to form ionic hydrides M2+ (H-)2
Group II M(s) + H2(g)
43
600C – 700C
MH2(s)
Group 2: Hydrides • Form hydrides of type MH2 • Be, Mg – Little tendency − − − −
Polymeric hydrides (BeH2 ) Three centre two electron bond BeH2 is covalent MgH2 is partially ionic
• Ca, Ba, Sr – ionic hydrides
44
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 45
Reaction with Air / Oxygen
46
Reaction with Air / Oxygen Group I Elements •
All alkali metals form more than one type of oxide on burning in air (except lithium)
Group II Elements
47
•
All alkaline earth metals react slowly with air to form oxides
•
On burning in air, they form both oxide and nitride
Reaction with Air / Oxygen : Group 1 Elements •
Three types of oxides:
normal oxides peroxides
superoxides Abundant supply
O2– oxide ion 48
1 O2 2
O22–
peroxide ion
O 2O2– 2
superoxide ion
Reaction with Air / Oxygen : Group 1 Elements Type of oxide formed depends on
1. supply of oxygen 2. reaction temperature
3. charge density of M+
49
Reaction with Air / Oxygen : Group 1 Elements •
Lithium when it is burnt in air, it forms normal oxide only 4Li(s) + O2(g)
50
180 C
2Li2O(s)
lithium oxide
Reaction with Air / Oxygen : Group 1 Elements •
Sodium
when it is burnt in an abundant supply of oxygen forms both the normal oxide and the peroxide
4Na(s) + O2(g) 2Na2O(s) + O2(g)
excess
51
180 C
2Na2O(s)
sodium oxide 300 C
2Na2O2(s)
sodium peroxide
Reaction with Air / Oxygen : Group 1 Elements •
Potassium, rubidium and caesium form All three types of oxides when burnt in sufficient supply of oxygen
52
Reaction with Air / Oxygen : Group 1 Elements Group I element
Normal oxide
Peroxide
Superoxide
Li
Li2O
–
–
Na
Na2O
Na2O2
–
K
K2O
K2O2
KO2
Rb
Rb2O
Rb2O2
Cs
Cs2O
Cs2O2
RbO2 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 53
Reaction with Air / Oxygen : Group 1 Elements
The electron cloud of the superoxide ion is greatly distorted by the small lithium ion 54
Reaction with Air / Oxygen : Group 1 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
Super oxides are generally bright coloured
They exhibit paramagnetic character due to unpaired electron 55
Reaction with Air / Oxygen : Group 1 Elements KO2 used as oxygen generators and CO2 scrubbers in spacecrafts and submarines 4KO2 + 2H2O 4KOH + 3O2 2KOH + CO2 K2CO3 + H2O
56
Reaction with Air / Oxygen : Group 2 Elements Supero
Group II element
Normal oxide
Peroxide
Be
BeO
–
–
Mg
MgO
–
–
Ca
CaO
–
–
Sr
SrO
-
–
Ba
BaO
-
–
xide
All these oxides are basic in nature (except beryllium oxide which is amphoteric) 57
Solubility • ΔG0 = ΔH0 - TΔS0
• General rule: • Compounds that contain widely differing radii are soluble in water • Difference in size favours solubility (>80pm)
Thermodynamics of dissolution • Entropy favours dissolution • Hydration energy of a smaller ion is larger • ΔLH = 1 / (r+ + r- )
and
ΔHydH = (1 / r+ ) + (1 / r- )
• Ion size assymmetry results in exothermic dissolution
• If both are small, both ΔLH and ΔHydH may be large, but enthalpy of dissolution may not be very exothermic 58
Solubility • The solubility of compounds increases with increase in ionic size of metal • Fluorides, oxides, hydroxides
• The solubility of compounds decreases with increase in ionic size of metal • Carbonates, sulphates, nitrates, halides (except fluorides)
59
Processes involved in Dissolution and their Energetics
•
60
Two processes are 1.
Breakdown of the ionic lattice
2.
Hydration
Hsolution
Na+(aq)
NaCl(s)
Na+(g)
ΔH
o solution
+
Cl-(aq)
Cl -(g)
ΔH
o hydration
ΔH
= (-772 +776) kJ mol1 61
+
= +4 kJ mol1
o lattice
o 0 , we expect the solids to dissolve in If ΔHsolution water
Solubility as
o ΔHsolution becomes more –ve (less +ve)
o ΔHsolution Solids (e.g. NaCl) with small +ve values are also soluble in water if the dissolution involves an increase in the entropy of the system.
ΔG
o solution
62
ΔH
o solution
TS
o solution
ΔG
o solution
ΔG
o solution
ΔH
o solution
0
TS
o solution
Spontaneous dissolution
TS
o is always positive solution
Dissolution with slightly positive can be spontaneous
63
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,
64
o L and
H
o hydration
H
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
ΔH
o hydration
less –ve down the group
ΔH
o lattice
+ve constant
Solubility down the group 65
Trends and Interpretations
rSO 2 r 4
ΔH o L
1
constant
rSO 2 r 4
ΔH
o solution
less –ve down the group 66
(-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,
67
o L and
H
o hydration
H
become less -ve
o L more rapidly than
H
o hydration
Trends and Interpretations
ΔH
o solution
more –ve down the group
(-ve)
ΔH
o hydration
less rapidly down the group
(+ve)
ΔH
o lattice
more rapidly down the group
less +ve down the group
Solubility down the group
68
General Rules 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
69
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
70
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
71
Thermal Stability • ΔG0 = ΔH0 - TΔS0
• The ΔG0 for the decomposition of a solid becomes negative when TΔS0 > ΔH0 • ΔH0 depends on
(example carbonates)
= Enthalpy of decomposition + (Lattice Enthalpy of Carbonate - Lattice Enthalpy of Oxide) • Enthalpy of decomposition is generally large and positive
• Metals having small cations, increases the lattice enthalpy of oxide more than that of the carbonate / sulphate / hydroxide / peroxide • Therefore, Lattice enthalpy plays an important role in deciding the stability. 72
Thermal decomposition reactions Metal carbonates
M2CO3(s) MCO3(s)
heat heat
M2O(s) + CO2 MO(s) + CO2
Metal hydroxides 2MOH(s) M(OH)2(s) 73
heat heat
M2O(s) + H2O(g) MO(s) + H2O
Relative thermal stability can be measured in two ways A higher decomposition temperature a greater thermal stability BeCO3(s) MgCO3(s)
CaCO3(s) SrCO3(s) BaCO3(s) 74
100C
BeO(s) + CO2(g) 540C
MgO(s) + CO2(g) 900C
CaO(s) + CO2(g) 1290C
SrO(s) + CO2(g) 1360C
BaO(s) + CO2(g)
Relative thermal stability can be measured in two ways 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
75
a thermally more stable compound
Be(OH)2(s)
Mg(OH)2(s)
MgO(s) + H2O(g)
Ca(OH)2(s)
Sr(OH)2(s)
SrO(s) + H2O(g)
Ba(OH)2(s)
BeO(s) + H2O(g) H = +54 kJ mol–1 H = +81 kJ mol–1
CaO(s) + H2O(g) H = +109 kJ mol–1 H = +127 kJ mol–1
BaO(s) + H2O(g) H = +146 kJ mol–1
Factors affecting thermal stability 1.
Polarizing power of cation
2.
Polarizability of polyatomic anion
3.
Lattice enthalpy of metal oxide produced
76
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
Polarizability as the size of anion 77
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
78
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 79
Interpretation of trends in thermal stability of carbonates and hydroxides 2.
Thermal stability down the groups
∵ size of cations down the groups ∴ (a) groups (b)
80
charge density/polarizing power of cation down the lattice enthalpies of MO/M2O down the groups
MgCO3(s) more polarized
BaCO3(s) less polarized
more favourable heat
less favourable heat
MgO(s) + CO2(g) more stable
BaO(s) + CO2(g) less stable
Thermal stability of carbonates down the groups
81
Effect of sizes of the cations on thermal stability of the carbonates and hydroxides of both Groups I and II metals 82
Reaction with Water
83
Action of Water • Both alkali and alkaline metals react with water • Respective Hydroxides and Hydrogen gas are formed
• Reactivity increases down the group • Type : Slow to explosive reactions • Na
84
K
Reactions with water or steam Group I
2M(s) + H2O(l)
heat
2MOH(aq) + H2(g)
Group II M(s) + 2H2O(l)
heat
M(OH)2(aq) + H2(g)
Mg reacts with steam but not cold water Mg(s) + H2O(g)
heat
MgO(s) + H2(g)
Be has no reaction with either water or steam
Reaction with ammonia • Exhibited both by Group I and II metals
• All show Blue colour • Ammoniated electron is present in these solutions, as the electron is solvated by ammonia
• Intensity of blue color increases with metal concentrations • High electrical conductivity • This solution show Magnetic properties
• Reducing property of solution of metal in ammonia (selective reducing action in organic chemistry) • These solution scan be used to prepare any desired oxide, by ing calculated quantities of oxygen gas through the solutions 86
Hydroxides
87
Group 1 Hydroxides of type MOH • These hydroxides are Strong bases • Basic strength / basic character / solubility in water / thermal stability − LiOH < NaOH < KOH < RbOH < CsOH • LiOH decomposes on heating to give water and Li2O
88
Group 2: Hydroxides of type M(OH)2 • All group 2 metals form hydroxides • Reaction of oxides with water gives hydroxides • Be(OH)2 Amphoteric
Mg(OH)2
Ca(OH)2 , Sr(OH)2 , Ba(OH)2
Weakly Basic
Strongly Basic
• Weaker bases than alkali metal hydroxides − Higher IE, smaller ionic size, higher charge on metal ion.
89
Group 2: Hydroxides • The solubility of the hydroxides in water increases with
increase in atomic number of the cation. • Be(OH)2 < Mg(OH)2 < Ca(OH)2 < Sr(OH)2 < Ba (OH)2 insoluble
insoluble
sp. soluble
soluble
soluble
• The solubility of hydroxides depend mainly on two facts.
− The lattice energy required to dissociate the components of hydroxide. This decreases from beryllium to barium. − The hydration energy of cation M2+. This decreases from beryllium to barium as the size of cation increases. • Both lattice and hydration energies decrease down the group, the decrease in lattice energy is more rapid than the hydration energy and so their 90 solubility increases on descending the group.
Compounds
Solubility / mol per 100 of water
Mg(OH)2
0.02 103
Ca(OH)2
1.5 103
Sr(OH)2
3.4 103
Ba(OH)2
15 103
Compounds
Solubility / mol per 100 of water
MgSO4
1800 104
CaSO4
11
104
SrSO4
0.71 104
BaSO4
0.009 104
down the group
down the group
Compounds
Solubility / mol per 100 of water
Mg(OH)2
0.02 103
Ca(OH)2
1.5 103
Sr(OH)2
3.4 103
Ba(OH)2
15 103
In general,
Size and/or charge of the anion Polarizability of anion Covalent character
Compounds
Solubility / mol per 100 of water
MgSO4
1800 104
CaSO4
11 104
SrSO4
0.71 104
BaSO4
0.009 104
Solubility in water
Group I: Carbonates • Type M2CO3 • Solubility in Water: • Increases as the size (atomic number) of cation increases. • Li2CO3 < Na2CO3 < K2CO3 < Rb2CO3 < Cs2CO3 • Low
--------------High-------------------------------
• Thermal Stability • Li2CO3 < Na2CO3 < K2CO3 < Rb2CO3 < Cs2CO3 • Low 93
--------------High-------------------------------
Group II: Carbonates • All form carbonates of the type MCO3 • Solubility in Water: • Insoluble in neutral medium, soluble in acidic medium
• Solubility decreases down the group • BeCO3 > MgCO3 > CaCO3 > SrCO3 > BaCO3 • Carbonates are more soluble in a solution containing CO2 Bicarbonates
• All carbonate solutions undergo the above reaction • Bicarbonates cannot be obtained in solid form but are known in solution state only. • Na, K, Rb, Cs bicarbonates are the only ones that can be obtained in solid 94 state.
Group II: Carbonates • Thermal Stability
Increases down the group
• BeCO3 < MgCO3 < CaCO3 < SrCO3 < BaCO3 • BeCO3 must be stored under CO2
95
Reaction with Nitrogen
96
Alkali Metal : Reaction with Nitrogen family • Lithium forms Nitrides (exceptions w.r.t alkali metal reactions) • Other metals form Azides (MN3) • They form binary compounds with other family of N
• The binary compounds undergo hydrolysis in water to form ammmonia, phosphine, asine, stibine etc 97
Alkaline Earths : Reaction with nitrogen family • All metals form nitrides – M3N2 • Ease of formation of nitrides decreases down the group
• These nitrides are stable up to 10000C • Get hydrolysed in water to give ammonia
98
Halides
99
Group 1 : Halides • MX
2M(s) + Cl2(g)
heat
2MCl(s)
• Ionic compounds , high lattice energies, • Stability
• The order of enthalpy of formation of a metal halide is
• Fluoride > Chloride > Bromide > Iodide • Fluorides are highly stable
100
Group 1: Halides • Trends in melting and boiling points of halides: • For a given alkali metal, the melting points and boiling points:
• Fluoride > Chloride > Bromide > Iodide • For a given halogen, the melting and boiling points
• Lithium
< Sodium > Potassium > Rubidium > Caesium
− Due to covalent character of Li compound
101
Halides : Ionic Character • The order of ionic character is • LiX < NaX < KX < RbX < CsX
• MF > MCI > MBr > MI − (same metal, different halogen)
102
Group 2: Halides • MX2
M(s) + Cl2(g)
heat
MCl2(s)
• When crystallized from solutions they form hydrated salts • Anhydrous CaCl2, SrCl2 and BaCl2 can be prepared by heating the hydrated salts.
103
Group 2: Halides • Alkaline earth metals combine with halogen on heating to form MX2 type salts. • Be Halides are covalent • Other halides are ionic
• Ionic Character − BeX2 < MgX2 < CaX2 < SrX2 < BaX2 − MI2 < MBr2 < MCl2 < MF2
104
Structure of BeCl2 molecules
A) At high temperatures, BeCl2 occurs as a gaseous molecule with only four electrons around Be. B) In the solid state, BeCl2 occurs in long chains with each Cl bridging two 105 atoms, which gives each Be an octet. Be
Group 2: Halides • Except BeCl2, all other halides are hygroscopic • Extent of hydration decreases down the group − Be and Mg halides hydrolyse on heating − Ca, Sr, Ba halides get dehydrated on heating − Calcium chloride has a strong affinity for water
• Solubility order • Fluorides are readily soluble
106
− BeF2 > MgF2 > CaF2 < SrF2 < BaF2 − BeX2 > MgX2 > CaX2 > SrX2 > BaX2 − MF2 < MCl2 < MBr2 < MI2
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 107
Group II Sulfates • Obtained by action of dil sulfuring acid on − − − −
Metal Metal oxide Metal hydroxide Carbonate
• Sulfates of Be, Mg, Ca crystalise as Hydrated salts • BeSO4 . 4H2O MgSO4 . 7H2O
CaSO4 . 2H2O
• Sulfates of Sr and Ba crystallise without water of crystallisation
108
Group II Sulfates • Solubility in water •
BeSO4 > MgSO4 > CaSO4 >
•
fairly soluble
SrSO4 > BaSO4
completely insoluble
• Thermal stability increases down the group
•
109
BeSO4 < MgSO4 < CaSO4 < SrSO4 < BaSO4
Nitrates • Group 1 : All form nitrates of type MNO3 • Group 2 : All form nitrates of type M(NO)2 • All are ionic • All are soluble in water
110
General Reactions of Alkali Metals
111
Important Reactions of the Alkaline Earth Metals - I
The metals reduce O2 to form the oxide: 2 M(s) + O2 (g)
MO(s)
Barium also forms the peroxide BaO (s).
The Metals of higher atomic weight reduce water to form hydrogen gas: M(s) + 2 H2O(l) M(OH)2 (aq) + H2 (g) M = Ca, Sr and Ba Be and Mg form an adherent oxide coating that allows only slight reaction.
112
Important Reactions of the Alkaline Earth Metals - I
The metals reduce halogens to form ionic halides:
M(s) + X2(-)
MX2 (s)
X = F, Cl, Br, I
Most of the metals reduce hydrogen to form ionic hydrides.
M(s) + H2 (g) MH2 (s)
113
all except Be
Important Reactions of the Alkaline Earth Metals - II Most of the metals reduce nitrogen to form ionic nitrides:
3 M(s) + N2 (g)
M3N2 (s)
all except Be
Except for amphoteric BeO, the oxides are basic:
MO(s) + H2O(l) M(OH)2 (aq) All carbonates undergo thermal decomposition to the oxide: MCO3 (s) MO + CO2 This reaction is used to produce CaO (lime) in huge amounts from naturally occurring limestone, and was the reaction used to generate carbon dioxide to smother the graphite fire in the Chernobyl reactor.
114
General Reactions of Alkaline Earth Metals
115
Diagonal relationship
116
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
117
118
Similarities of Be and Al •
Be and Al have the same electronegativity (Be I.0 and AI 1.5) and their charge/radius ratios are very
•
The standard oxidation potentials of both Be and are of nearly the same order (Be = 1.97V; Al= l.7V)
•
Since the polarizing power of both Be and Al are nearly the same, the covalent character of their compounds also similar.
•
Both Be and Al are rendered ive on treatment with conc. HNO 3.
•
Unlike alkaline earth metals Be does not get readily attacked by dry air. (like Al)
•
Both Be and Al reacts very slowly with dilute mineral acids due to the presence of
oxide layer. •
Both Be and Al react with alkalis liberating H2.
•
Both Be and AI form carbides which on hydrolrolysis liberate methane.
•
Both form nitrides when heated in nitrogen which give ammonia by the reaction with water.
•
Both form oxides which are amphoteric.
119
Halides of both Be and AI contain halogen bridge bonds
•
120
121
Compounds of Na
122
Na2CO3 – Solvay Process
123
124 C a r b o n a t i n g To w e r
125
Sodium Carbonate
126
Caustic Soda – NaOH - Nelson Cell
127
NaOH by Castner Kellner Cell
128
NaOH by Castner Kellner Cell • Castner-Kellner method also known as Mercury Cathode Method. • In this method, the electrolytic cell contains three compartments. (i) Mercury in the outer compartment acts as a cathode while in middle compartment acts as an anode due to induction. (ii) Graphite rods in the outer compartments acts as anode while the iron rods in the middle compartment acts as a cathode.
• Sodium liberated at mercury cathode in the out compartments dissolve in mercury forming sodium amalgam which moves into middle compartment where it react with water at cathode forming NaOH, H2 and Hg. Cl2 gas is liberated at graphite anodes in the outer compartment.
129
Sodium Hydroxide
130
Sodium Hydroxide
131
Sodium Hydroxide
Sodium Beryllate Sodium Aluminate Sodium Stannite Sodium Plumbite Sodium Zincate
132
Sodium Hydroxide
133
Sodium Sulphate - Salt Cake anhydrous Na2SO4 - Glauber’s Salt Na2SO4 .10 H2O •
Preparation
•
The salt cake (anhydrous sodium sulphate) is dissolved in water and the solution is subjected to crystallization.
•
Above 32 °C the anhydrous salt separates.
•
Below 32 °C, the decahydrate salt crystallises out from the aqueous solution.
•
Saturated solution of the decahydrate, on cooling below 12 °C, gives crystals of heptahydrate.
•
Properties
Pb ( NO3 ) 2 Na2 SO4 PbSO4 2 NaNO3 BaCl 2 Na2 SO4 BaSO4 2 NaCl Sr ( NO3 ) 2 Na2 SO4 SrSO4 2 NaNO3
• Uses: It is used in textile industry, medicines as purgative, manufacture of glass plates and sodium salts. 134
Sodium Bicarbonate, Baking Soda, Na2HCO3 • Preparation By ing CO2 through Sodium Carbonate solution but industrially it is manufactured by Solvay's process. • Properties • It is sparingly soluble in Water • Solution is alkaline in nature • Uses
• On heating, it decomposes to give sodium carbonate. • The metal salts which gives basic metal carbonate with sodium carbonate gives normal carbonates.
• Sodium bicarbonate is used, as an antacid in medicine, in dry fire extinguishers, in baking powders and as mild antiseptic for skin infections. 135
Compounds of Alkaline Earth Metals
136
Magneisum Oxide - MgO – Magnesia • Preparation:
1. Calcination of Magnesite (MgCO3) MgCO3 MgO + CO2 2. Heating Mg(NO3)2 or Mg(OH)2
Mg(NO3)2 MgO + 4 NO2 + O2 Mg(OH)2 MgO + H2O
• Properties: • Light infusible white solid • They have high MP 3073K
• Used as refractory material due to the above property 137
Magneisum Oxide - MgO – Magnesia • Chemical properties • Hydrolyses in water to form insoluble Mg(OH)2 2MgO H 2 O 2Mg(OH) 2 • Being basic, reacts with acids to form respective salts MgO 2HCl MgCl 2 H 2 O • Gives Mg on reduction with Carbon at high tempertaures MgO + C Mg + CO
MgO 3C MgC 2 CO • Uses: • Sorel’s cement – used in Dentistry
MgCl2. 5MgO.xH2O
• As an antacid • As an insulator when mixed with asbestos 138
Magnesium Hydroxide Mg(OH)2 • Preparation: 1. By the hydrolysis of MgO 2. By treating MgCl2 with Ca(OH)2 MgCl2 + Ca(OH)2 Mg(OH)2 + CaCl2 • White powdery substance • Sparingly soluble in water
• Used as an Antacid under the name Milk of Magnesia
139
Magnesium Carbonate MgCO3 • Preparation: Hot Magnesium sulfate with sodium bicarbonate
MgSO4 2 NaHCO3 MgCO3 Na2 SO4 H 2O CO2 • Basic magnesium carbonate H 2O
MgSO4 Na2CO3 MgCO3 .Mg (OH ) 2 Na2 SO4 CO2 (Basic Magnesium Carbonate / Magnesia alva)
MgCO3 .Mg (OH ) 2 3CO2 H 2O Mg ( HCO3)2
Mg ( HCO3)2 MgCO3 H 2O CO2
• A solution containing 12% MgCO3 per 100 cc of water containing dissolved CO2 in called Fluid Magnesia 140
Magnesium Sulphate MgSO4.7 H2O Epsom Salt • Preparation • From Magnesite : heating with dil. Sulfuric acid
MgCO3 H 2 SO4 MgSO4 CO2 H 2O • From Dolomite MgCO3.CaCO3 2H 2 SO4 MgSO4 CaSO4 2CO2 2H 2O • From Keiserite (commercial method). Boil with water and cool
MgSO4.H 2O H 2O MgSO4.7 H 2O • Properties • Colourless efflorescent solid C C C MgSO 4 .7H 2 O 30 MgSO 4 .6H 2 O 150 MgSO 4 .H 2 O 200 MgSO 4
141
Magnesium Sulphate MgSO4.7 H2O Epsom Salt • On heating, it decomposes C 4MgSO 4 250 4MgO 2SO 2 2SO 3 O 2
• On heating with Carbon, it gets reduced
2MgSO 4 C 2MgO 2SO 2 CO 2 • It forms double salts with alkali metal sulfates
K 2SO 4 .MgSO 4 .6H 2 O
142
Calcium Oxide / Quick Lime / Burnt Lime • Preparation: Decomposition of Limestone •
C CaCO 3 900 CaO CO 2
• Reacts with water with a hissing noise to form Slaked lime Ca (OH)2 (Rxn is known as Slaking of lime) ΔH = -15 kcal/mol
CaO H 2 O Ca (OH) 2 • Milk of lime : paste of lime in water • Lime water : Clear filtrate • “Limelight” in oxy hydrogen flame
143
Calcium Oxide / Quick Lime / Burnt Lime • CaO gives Calcium Silicate with silica and Calcium Phosphate with P4O10
CaO SiO2 CaSiO3 6CaO 2 P2O5 2Ca3 ( PO4 ) 2 • Forms Calcium Carbide on heating with carbon (2000 deg C) C CaO 3C 2000 CaC 2 CO
• Calcium Carbide + water gives Calcium Cyanamide CaC 2
C N 2 1000
CaCN 2
C
Calcium cyanamide
−
144
Calcium Cyanamide + C = Nitrolim - a Fertiliser
Calcium Oxide / Quick Lime / Burnt Lime • Uses
• It is used as a drying agent. • It is used in the manufacture of bleaching powder. • It is used in the manufacture of calcium. carbide, cement, glass, lime mortar, etc. • It is used in the purification of sugar.
145
Calcium Hydroxide / Slaked Lime / Milk of Lime • Preparation: Slaking of lime • Properties • White amorphous solid • Sparingly soluble in water •
On heating, loses water molecule to form Lime CaO
• Action of CO2
Ca (OH) 2 CO 2 CaCO 3 H 2 O CaCO 3 H 2 O CO 2 Ca (HCO 3 ) 2 Insoluble
So luble
• Similar reaction with SO2 gas is seen when Calcium bisulphite is 146 formed
Calcium Hydroxide / Slaked Lime / Milk of Lime • Reaction with Ammonia Ca (OH) 2 2NH 4 Cl Heat CaCl 2 2NH 3 H 2 O
• Reaction with Chlorine 3Ca(OH)2 + 2Cl2
below 35oC
slaked lime
Ca(OCl)2.Ca(OH)2.CaCl2. 2 H2 O +
H2 O
bleaching powder
•
Bleaching powder is a calcium salt of hypochlorous acid (HOCl)
•
Ca(OCl)2
•
Uses: It is used
1. for absorbing acid gases. 2. in the manufacture of bleaching powder and caustic soda. 3. in the production of lime mortar for construction of buildings, whitewashing buildings 4. in glass making, tanning industry and for purification of sugar 5. for the preparation of NH3 from NH4Cl in Solvay process
147
6. as lime water in laboratories
Gypsum CaSO4.2H2O • Preparation
CaCl 2 H 2SO 4 CaSO 4 2HCl CaCl 2 Na 2SO 4 CaSO 4 2NaCl • Properties • White crystalline solid • Solubility decreases on increase in temperature • Action of heat – Calcium sulphate hemihydrate (plaster of paris) is formed C 2[CaSO 4 .2H 2 O] 120 (CaSO 4 ) 2 .H 2 O 3H 2 O
Plaster of Paris C [(CaSO 4 ) 2 .H 2 O] 200 2CaSO 4 H 2 O
Dead burnt plaster 148
heated 2CaSO 4 Strongly 2CaO 2SO 2 O 2
Calcium Carbonate, CaCO3 • Naturally found as limestone, marble, chalk
• Preparation
Ca (OH) 2 CO 2 CaCO 3 H 2 O
CaCl 2 Na 2 CO 3 CaCO 3 2NaCl • White fluffy powder insoluble in water. But dissolves in water in the presence of carbon-di-oxide to form calcium bicarbonate
CaCO 3 H 2 O CO 2 Ca (HCO 3 ) 2
149
Mortar • It is also known as lime mortar. • It is an intimate mixture of 1 part of slaked lime, 3 parts of sand and water made into paste. • This is used to bind the bricks firmly. • Setting of mortar involves the following steps. • (i) Mortar loses water on of evaporation. • (ii) Carbon dioxide is absorbed from the air converting into calcium carbonate which acts as a binding material. • (iii) Slaked lime reacts with silica forming calcium silicate which gives hardness.
150
Cement • The name portland cement was given to it by Joseph Aspidin (a mason!) because when it is mixed with sand and water it hardens like the lime stone querried at Portland in England. • Composition : CaO 50 to 60 %; SiO2: 20 to 25%; Al2O3 : 5 to 10 %; MgO :2 to 3%; Fe2O3 I to 2% and SO3 1 to 2%. • If lime is excess the cement cracks during setting but if it is less the cement will be weak. • Excess of Al2O3 will make cement quick drying • The raw materials for the manufacture of cement are limestone and alumino silicates (clay, sand and shales). When the powdered raw materials are heated in a rotary kiln, sintered clinker will be obtained. • 151
The setting of cement by mixing with water is due to hydration of the molecules and their rearrangement.
S - Block Metals in Biological Systems
152
Biological functions of Sodium and Potassium ions •
Sodium and potassium are the most common cations in biological fluids.
•
Sodium ion is the major cation of extracellular fluids of animals and in blood plasma, including human beings which is known to activate certain enzymes in the animal body
•
These ions participate in the transmission of nerve signals.
•
They also regulate flow of water across cell membranes and in transport of sugars, amino acids into the cells.
•
Potassium ions are the most abundant cations within cell fluids, where they activate many enzymes that participate in oxidation of glucose to produce adenosine triphosphate (ATP).
•
A typical 70 kg adult contains about 90 g of Na+ ions and 170 g of K+ ions.
•
The daily requirement of sodium and potassium is about 2 g each.
153
Biological functions of Magnesium and Calcium •
Magnesium is an important constituent of chlorophyll.
•
Mg2+ and Ca2+ ions are also responsible for the transmission of electrical impulses along the nerve fibre and the contraction of muscles
•
Calcium ions are essential for the formation of bones and teeth
•
It also plays important roles in maintaining rhythm of heart, clotting of blood, neuromuscular function, interneuronal transmission, cell membrane integrity, etc.
•
The calcium concentration in plasma is regulated at about 100 mg L-1. It is maintained by two hormones, calcitonin and parathyroid hormone.
•
The substance present in bones is continuously solubilized and redeposited to the extent of 400 mg per day in man.
•
All this calcium es through the plasma.
154
The END
The s-block elements
s-Block Elements • Similarities • Highly reactive metals • Strong reducing agents • Form ionic compounds • Fixed oxidation state • Group I : +1 • Group II : +2
Variation in Physical Properties of s-block Elements 1. Atomic Radius and Ionic Radius 2. Ionization Enthalpies 3. Hydration Enthalpies 4. Melting Points 5. Electronegativity
4
Atomic and Ionic Radii • The atoms and ions of alkali metals are largest in their corresponding periods. • Atomic size
Li < Na < K < Rb < Cs
• Ionic Radius
Li+ < Na + < K + < Rb + < Cs +
• Atomic volume
Li < Na < K < Rb < Cs
• Charge density
Li > Na > K > Rb > Cs
5
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
down the groups ∵ the outermost electrons are further away from the nuclei Group II < Group I ∵ ENC from left to right across the periods
6
On moving down the groups, first sharply (e.g. from Li to K)
then slowly (e.g. from K to Fr) There is a sharp in NC from 19K to 37Rb Outermost e is drawn closer to the nucleus
The inner d-electrons (of Rb, Cs, Sr, Ba) have poor shielding effect on the outermost electrons transition contraction 7
Ionisation Enthalpy
8
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
Both atomic radius and ENC down the groups Atomic radius is more important IE down the groups
9
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
For Group I elements, 2nd IE >> 1st IE because the 2nd electron is closer to the nucleus and is poorly shielded by other electrons in the same shell which is completely filled.
10
For Group II elements, 3rd IE >> 2nd IE Similar reasons can be applied
Ionization Enthalpy
Variations in the first and second ionization enthalpies of Group I elements
11
Ionization Enthalpy
Variations in the first, second and third ionization enthalpies of Group II elements
12
Electronegativity
13
Electronegativity • Relatively LOW Electronegativity • These metals have more tendency to lose electron rather than to gain an electron.
• The electronegativity values decreases down the group from Li to Cs
Li > Na > K > Rb > Cs
14
ELECTRONEGATIVITY 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
All have low electronegativity => Electropositive EN down the group EN : Group II > Group I (∵ greater ENC)
Hydration
16
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
17
always has a negative value
Hydration energies • Group I • Alkali metal ions are highly hydrated. • The smaller the ionic size, the higher the degree of hydration. • Primary and secondary shell of hydration • Li ion is very small, it is heavily hydrated. • Li ion is tetrahedrally surrounded by four water molecules using its four sp3
hybrid • Group 2
• They have higher hydration energies than Alkali metals due to smaller sizes 18
Hydration energies • In aqueous solutions, degree of hydration decreases from Li+ to Cs+ due to increase in size • Ionic radii of hydrated alkali metal ions also decreases from Li+ to Cs+
• Formation of hydrated salts : − Li > Na > K Salts. − Rb and Cs salts are not hydrated • Ionic Mobility : Cs+ > Rb+ > K + > Na + > Li +
19
Hydration energies Hydration Group I ion
Enthalpy (kJ mol–1)
Hydration Group II ion
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
20
Group II > Group I ∵ Group II ions have higher charge and small size higher charge density stronger ion-dipole interaction
Melting & Boiling 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
21
Group I
Melting
Group II
Melting
element
Point (C)
element
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
down the groups ∵ ionic radii down the groups charge density interaction between ions and electron sea 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 22
STRUCTURE Group I
Densities of Li, Na, K are lesser than that of water
Densit y (g cm3)
Group II
Densit y (g cm3)
Li
0.53
Be
1.86
Na
0.97
Mg
1.74
K
0.86
Ca
1.55
Rb
1.53
Sr
2.54
Cs
1.90
Ba
3.59
Fr
-
Ra
-
Density • Alkali metals have low density. − The reason this is that they have large atomic sizes.
• Density gradually increases on moving down the group from Li to Cs Li < Na < K < Rb < Cs.
• Anomaly: K is lighter than Na
24
Effect of light • Alkali metals when irradiated with light emit electrons with ease due to low ionization enthalpies. • This phenomenon is used in photoelectric cells, particularly caesium and potassium are used as electrodes in photoelectric cells.
25
Flame Colouration
26
Flame Colouration Most s-block elements and their compounds give a characteristic flame colour in the flame test Group I element Li Na K Rb Cs
27
Flame colour Crimson Golden yellow Lilac Bluish red Blue
Flame Colouration Most s-block elements and their compounds give a characteristic flame colour in the flame test Group II element Be Mg Ca Sr Ba
28
Flame colour Brick red Blood red Apple green
Ca, Sr, Ba
Beryllium and magnesium atoms are smaller and their electrons being strongly bound to the nucleus are not excited to higher-energy levels.
Flame Colouration 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)* 29
[Ne] 3p1
cool
Na(g) [Ne] 3s1
+ golden yellow light Visible region
Flame Colouration Mechanism : For salts of s-block elements, the metal ions of the salts are first converted to metal atoms Na2CO3(s) Na+Cl
Na(g) 30
Na(g)*
Conc. HCl
heat heat cool
Na+Cl (more volatile)
Na(g) + Cl(g)
Na(g)* Na(g)
+ golden yellow light
Complex Formation
31
Complex formation • In order to form complex compounds, a metal must possess the following characteristics. − Small size − High effective nuclear charge − Tendency to accept electrons (i.e., presence of vacant orbitals) • Since alkali metals have none of these characteristics they have little tendency to form complexes. • Lithium and Beryllium forms certain complexes. (Due to their small sizes) • The complex forming tendency fall markedly down the groups as the atomic size increases
32
Complex formation : Weak Tendency 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
33
Complex formation : Weak Tendency Reasons
2. s-block cations (M+, M2+) have relatively low charge densities less polarizing and less able to accept lone pairs from ligands.
34
Complex formation : Owing to its high charge density, Be2+ can form complexes
35
Electropositive Character • The electropositive character increases down the group from Li to Cs because ionization enthalpy
decreases down the group Li > Na > K > Rb > Cs.
Eo 36
Group I
(V)
Group II
(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) down the groups Group I > Group II
Reducing Property • Powerful reducing agents • Li > Na < K = Rb > Cs (E0) • Reasons − Heat of sublimation − Ionisation enthalpy − Hydration energy
37
Reaction with Hydrogen
38
Reaction with Hydrogen Alkali metals react with hydrogen to form ionic hydrides M+H-. The reaction of alkali metals with hydrogen decreases from Li to Cs
Group I 2M(s) + H2(g)
300C – 500C
2MH(s)
Group 1 : Hydrides • The order and reactivity with hydrogen • Li > Na > K > Rb > Cs • The ionic character of the bonds in these hydrides Increases from LiH to CsH • LiH < NaH < KH < RbH < CsH
• Stability • LiH > NaH > KH > RbH > CsH
40
LAH
Dry ether
4LiH + AlCl3
LiAlH4 + 3LiCl
• Powerful reducing agent • Tetrahedral • Selective reducing agent − Reduces carbonyl compounds to alcohols.
• It reacts violently with water, so it is necessary to use absolutely dry organic solvents • Also reduces several inorganic substances
41
Sodium tetrahydridoborate (sodium borohydride)
• NaBH4 • Can be used even in aqueous solutions • Na and K hydrides are useful
42
Reaction with Hydrogen
Alkaline earth metals react with hydrogen to form ionic hydrides M2+ (H-)2
Group II M(s) + H2(g)
43
600C – 700C
MH2(s)
Group 2: Hydrides • Form hydrides of type MH2 • Be, Mg – Little tendency − − − −
Polymeric hydrides (BeH2 ) Three centre two electron bond BeH2 is covalent MgH2 is partially ionic
• Ca, Ba, Sr – ionic hydrides
44
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 45
Reaction with Air / Oxygen
46
Reaction with Air / Oxygen Group I Elements •
All alkali metals form more than one type of oxide on burning in air (except lithium)
Group II Elements
47
•
All alkaline earth metals react slowly with air to form oxides
•
On burning in air, they form both oxide and nitride
Reaction with Air / Oxygen : Group 1 Elements •
Three types of oxides:
normal oxides peroxides
superoxides Abundant supply
O2– oxide ion 48
1 O2 2
O22–
peroxide ion
O 2O2– 2
superoxide ion
Reaction with Air / Oxygen : Group 1 Elements Type of oxide formed depends on
1. supply of oxygen 2. reaction temperature
3. charge density of M+
49
Reaction with Air / Oxygen : Group 1 Elements •
Lithium when it is burnt in air, it forms normal oxide only 4Li(s) + O2(g)
50
180 C
2Li2O(s)
lithium oxide
Reaction with Air / Oxygen : Group 1 Elements •
Sodium
when it is burnt in an abundant supply of oxygen forms both the normal oxide and the peroxide
4Na(s) + O2(g) 2Na2O(s) + O2(g)
excess
51
180 C
2Na2O(s)
sodium oxide 300 C
2Na2O2(s)
sodium peroxide
Reaction with Air / Oxygen : Group 1 Elements •
Potassium, rubidium and caesium form All three types of oxides when burnt in sufficient supply of oxygen
52
Reaction with Air / Oxygen : Group 1 Elements Group I element
Normal oxide
Peroxide
Superoxide
Li
Li2O
–
–
Na
Na2O
Na2O2
–
K
K2O
K2O2
KO2
Rb
Rb2O
Rb2O2
Cs
Cs2O
Cs2O2
RbO2 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 53
Reaction with Air / Oxygen : Group 1 Elements
The electron cloud of the superoxide ion is greatly distorted by the small lithium ion 54
Reaction with Air / Oxygen : Group 1 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
Super oxides are generally bright coloured
They exhibit paramagnetic character due to unpaired electron 55
Reaction with Air / Oxygen : Group 1 Elements KO2 used as oxygen generators and CO2 scrubbers in spacecrafts and submarines 4KO2 + 2H2O 4KOH + 3O2 2KOH + CO2 K2CO3 + H2O
56
Reaction with Air / Oxygen : Group 2 Elements Supero
Group II element
Normal oxide
Peroxide
Be
BeO
–
–
Mg
MgO
–
–
Ca
CaO
–
–
Sr
SrO
-
–
Ba
BaO
-
–
xide
All these oxides are basic in nature (except beryllium oxide which is amphoteric) 57
Solubility • ΔG0 = ΔH0 - TΔS0
• General rule: • Compounds that contain widely differing radii are soluble in water • Difference in size favours solubility (>80pm)
Thermodynamics of dissolution • Entropy favours dissolution • Hydration energy of a smaller ion is larger • ΔLH = 1 / (r+ + r- )
and
ΔHydH = (1 / r+ ) + (1 / r- )
• Ion size assymmetry results in exothermic dissolution
• If both are small, both ΔLH and ΔHydH may be large, but enthalpy of dissolution may not be very exothermic 58
Solubility • The solubility of compounds increases with increase in ionic size of metal • Fluorides, oxides, hydroxides
• The solubility of compounds decreases with increase in ionic size of metal • Carbonates, sulphates, nitrates, halides (except fluorides)
59
Processes involved in Dissolution and their Energetics
•
60
Two processes are 1.
Breakdown of the ionic lattice
2.
Hydration
Hsolution
Na+(aq)
NaCl(s)
Na+(g)
ΔH
o solution
+
Cl-(aq)
Cl -(g)
ΔH
o hydration
ΔH
= (-772 +776) kJ mol1 61
+
= +4 kJ mol1
o lattice
o 0 , we expect the solids to dissolve in If ΔHsolution water
Solubility as
o ΔHsolution becomes more –ve (less +ve)
o ΔHsolution Solids (e.g. NaCl) with small +ve values are also soluble in water if the dissolution involves an increase in the entropy of the system.
ΔG
o solution
62
ΔH
o solution
TS
o solution
ΔG
o solution
ΔG
o solution
ΔH
o solution
0
TS
o solution
Spontaneous dissolution
TS
o is always positive solution
Dissolution with slightly positive can be spontaneous
63
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,
64
o L and
H
o hydration
H
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
ΔH
o hydration
less –ve down the group
ΔH
o lattice
+ve constant
Solubility down the group 65
Trends and Interpretations
rSO 2 r 4
ΔH o L
1
constant
rSO 2 r 4
ΔH
o solution
less –ve down the group 66
(-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,
67
o L and
H
o hydration
H
become less -ve
o L more rapidly than
H
o hydration
Trends and Interpretations
ΔH
o solution
more –ve down the group
(-ve)
ΔH
o hydration
less rapidly down the group
(+ve)
ΔH
o lattice
more rapidly down the group
less +ve down the group
Solubility down the group
68
General Rules 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
69
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
70
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
71
Thermal Stability • ΔG0 = ΔH0 - TΔS0
• The ΔG0 for the decomposition of a solid becomes negative when TΔS0 > ΔH0 • ΔH0 depends on
(example carbonates)
= Enthalpy of decomposition + (Lattice Enthalpy of Carbonate - Lattice Enthalpy of Oxide) • Enthalpy of decomposition is generally large and positive
• Metals having small cations, increases the lattice enthalpy of oxide more than that of the carbonate / sulphate / hydroxide / peroxide • Therefore, Lattice enthalpy plays an important role in deciding the stability. 72
Thermal decomposition reactions Metal carbonates
M2CO3(s) MCO3(s)
heat heat
M2O(s) + CO2 MO(s) + CO2
Metal hydroxides 2MOH(s) M(OH)2(s) 73
heat heat
M2O(s) + H2O(g) MO(s) + H2O
Relative thermal stability can be measured in two ways A higher decomposition temperature a greater thermal stability BeCO3(s) MgCO3(s)
CaCO3(s) SrCO3(s) BaCO3(s) 74
100C
BeO(s) + CO2(g) 540C
MgO(s) + CO2(g) 900C
CaO(s) + CO2(g) 1290C
SrO(s) + CO2(g) 1360C
BaO(s) + CO2(g)
Relative thermal stability can be measured in two ways 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
75
a thermally more stable compound
Be(OH)2(s)
Mg(OH)2(s)
MgO(s) + H2O(g)
Ca(OH)2(s)
Sr(OH)2(s)
SrO(s) + H2O(g)
Ba(OH)2(s)
BeO(s) + H2O(g) H = +54 kJ mol–1 H = +81 kJ mol–1
CaO(s) + H2O(g) H = +109 kJ mol–1 H = +127 kJ mol–1
BaO(s) + H2O(g) H = +146 kJ mol–1
Factors affecting thermal stability 1.
Polarizing power of cation
2.
Polarizability of polyatomic anion
3.
Lattice enthalpy of metal oxide produced
76
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
Polarizability as the size of anion 77
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
78
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 79
Interpretation of trends in thermal stability of carbonates and hydroxides 2.
Thermal stability down the groups
∵ size of cations down the groups ∴ (a) groups (b)
80
charge density/polarizing power of cation down the lattice enthalpies of MO/M2O down the groups
MgCO3(s) more polarized
BaCO3(s) less polarized
more favourable heat
less favourable heat
MgO(s) + CO2(g) more stable
BaO(s) + CO2(g) less stable
Thermal stability of carbonates down the groups
81
Effect of sizes of the cations on thermal stability of the carbonates and hydroxides of both Groups I and II metals 82
Reaction with Water
83
Action of Water • Both alkali and alkaline metals react with water • Respective Hydroxides and Hydrogen gas are formed
• Reactivity increases down the group • Type : Slow to explosive reactions • Na
84
K
Reactions with water or steam Group I
2M(s) + H2O(l)
heat
2MOH(aq) + H2(g)
Group II M(s) + 2H2O(l)
heat
M(OH)2(aq) + H2(g)
Mg reacts with steam but not cold water Mg(s) + H2O(g)
heat
MgO(s) + H2(g)
Be has no reaction with either water or steam
Reaction with ammonia • Exhibited both by Group I and II metals
• All show Blue colour • Ammoniated electron is present in these solutions, as the electron is solvated by ammonia
• Intensity of blue color increases with metal concentrations • High electrical conductivity • This solution show Magnetic properties
• Reducing property of solution of metal in ammonia (selective reducing action in organic chemistry) • These solution scan be used to prepare any desired oxide, by ing calculated quantities of oxygen gas through the solutions 86
Hydroxides
87
Group 1 Hydroxides of type MOH • These hydroxides are Strong bases • Basic strength / basic character / solubility in water / thermal stability − LiOH < NaOH < KOH < RbOH < CsOH • LiOH decomposes on heating to give water and Li2O
88
Group 2: Hydroxides of type M(OH)2 • All group 2 metals form hydroxides • Reaction of oxides with water gives hydroxides • Be(OH)2 Amphoteric
Mg(OH)2
Ca(OH)2 , Sr(OH)2 , Ba(OH)2
Weakly Basic
Strongly Basic
• Weaker bases than alkali metal hydroxides − Higher IE, smaller ionic size, higher charge on metal ion.
89
Group 2: Hydroxides • The solubility of the hydroxides in water increases with
increase in atomic number of the cation. • Be(OH)2 < Mg(OH)2 < Ca(OH)2 < Sr(OH)2 < Ba (OH)2 insoluble
insoluble
sp. soluble
soluble
soluble
• The solubility of hydroxides depend mainly on two facts.
− The lattice energy required to dissociate the components of hydroxide. This decreases from beryllium to barium. − The hydration energy of cation M2+. This decreases from beryllium to barium as the size of cation increases. • Both lattice and hydration energies decrease down the group, the decrease in lattice energy is more rapid than the hydration energy and so their 90 solubility increases on descending the group.
Compounds
Solubility / mol per 100 of water
Mg(OH)2
0.02 103
Ca(OH)2
1.5 103
Sr(OH)2
3.4 103
Ba(OH)2
15 103
Compounds
Solubility / mol per 100 of water
MgSO4
1800 104
CaSO4
11
104
SrSO4
0.71 104
BaSO4
0.009 104
down the group
down the group
Compounds
Solubility / mol per 100 of water
Mg(OH)2
0.02 103
Ca(OH)2
1.5 103
Sr(OH)2
3.4 103
Ba(OH)2
15 103
In general,
Size and/or charge of the anion Polarizability of anion Covalent character
Compounds
Solubility / mol per 100 of water
MgSO4
1800 104
CaSO4
11 104
SrSO4
0.71 104
BaSO4
0.009 104
Solubility in water
Group I: Carbonates • Type M2CO3 • Solubility in Water: • Increases as the size (atomic number) of cation increases. • Li2CO3 < Na2CO3 < K2CO3 < Rb2CO3 < Cs2CO3 • Low
--------------High-------------------------------
• Thermal Stability • Li2CO3 < Na2CO3 < K2CO3 < Rb2CO3 < Cs2CO3 • Low 93
--------------High-------------------------------
Group II: Carbonates • All form carbonates of the type MCO3 • Solubility in Water: • Insoluble in neutral medium, soluble in acidic medium
• Solubility decreases down the group • BeCO3 > MgCO3 > CaCO3 > SrCO3 > BaCO3 • Carbonates are more soluble in a solution containing CO2 Bicarbonates
• All carbonate solutions undergo the above reaction • Bicarbonates cannot be obtained in solid form but are known in solution state only. • Na, K, Rb, Cs bicarbonates are the only ones that can be obtained in solid 94 state.
Group II: Carbonates • Thermal Stability
Increases down the group
• BeCO3 < MgCO3 < CaCO3 < SrCO3 < BaCO3 • BeCO3 must be stored under CO2
95
Reaction with Nitrogen
96
Alkali Metal : Reaction with Nitrogen family • Lithium forms Nitrides (exceptions w.r.t alkali metal reactions) • Other metals form Azides (MN3) • They form binary compounds with other family of N
• The binary compounds undergo hydrolysis in water to form ammmonia, phosphine, asine, stibine etc 97
Alkaline Earths : Reaction with nitrogen family • All metals form nitrides – M3N2 • Ease of formation of nitrides decreases down the group
• These nitrides are stable up to 10000C • Get hydrolysed in water to give ammonia
98
Halides
99
Group 1 : Halides • MX
2M(s) + Cl2(g)
heat
2MCl(s)
• Ionic compounds , high lattice energies, • Stability
• The order of enthalpy of formation of a metal halide is
• Fluoride > Chloride > Bromide > Iodide • Fluorides are highly stable
100
Group 1: Halides • Trends in melting and boiling points of halides: • For a given alkali metal, the melting points and boiling points:
• Fluoride > Chloride > Bromide > Iodide • For a given halogen, the melting and boiling points
• Lithium
< Sodium > Potassium > Rubidium > Caesium
− Due to covalent character of Li compound
101
Halides : Ionic Character • The order of ionic character is • LiX < NaX < KX < RbX < CsX
• MF > MCI > MBr > MI − (same metal, different halogen)
102
Group 2: Halides • MX2
M(s) + Cl2(g)
heat
MCl2(s)
• When crystallized from solutions they form hydrated salts • Anhydrous CaCl2, SrCl2 and BaCl2 can be prepared by heating the hydrated salts.
103
Group 2: Halides • Alkaline earth metals combine with halogen on heating to form MX2 type salts. • Be Halides are covalent • Other halides are ionic
• Ionic Character − BeX2 < MgX2 < CaX2 < SrX2 < BaX2 − MI2 < MBr2 < MCl2 < MF2
104
Structure of BeCl2 molecules
A) At high temperatures, BeCl2 occurs as a gaseous molecule with only four electrons around Be. B) In the solid state, BeCl2 occurs in long chains with each Cl bridging two 105 atoms, which gives each Be an octet. Be
Group 2: Halides • Except BeCl2, all other halides are hygroscopic • Extent of hydration decreases down the group − Be and Mg halides hydrolyse on heating − Ca, Sr, Ba halides get dehydrated on heating − Calcium chloride has a strong affinity for water
• Solubility order • Fluorides are readily soluble
106
− BeF2 > MgF2 > CaF2 < SrF2 < BaF2 − BeX2 > MgX2 > CaX2 > SrX2 > BaX2 − MF2 < MCl2 < MBr2 < MI2
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 107
Group II Sulfates • Obtained by action of dil sulfuring acid on − − − −
Metal Metal oxide Metal hydroxide Carbonate
• Sulfates of Be, Mg, Ca crystalise as Hydrated salts • BeSO4 . 4H2O MgSO4 . 7H2O
CaSO4 . 2H2O
• Sulfates of Sr and Ba crystallise without water of crystallisation
108
Group II Sulfates • Solubility in water •
BeSO4 > MgSO4 > CaSO4 >
•
fairly soluble
SrSO4 > BaSO4
completely insoluble
• Thermal stability increases down the group
•
109
BeSO4 < MgSO4 < CaSO4 < SrSO4 < BaSO4
Nitrates • Group 1 : All form nitrates of type MNO3 • Group 2 : All form nitrates of type M(NO)2 • All are ionic • All are soluble in water
110
General Reactions of Alkali Metals
111
Important Reactions of the Alkaline Earth Metals - I
The metals reduce O2 to form the oxide: 2 M(s) + O2 (g)
MO(s)
Barium also forms the peroxide BaO (s).
The Metals of higher atomic weight reduce water to form hydrogen gas: M(s) + 2 H2O(l) M(OH)2 (aq) + H2 (g) M = Ca, Sr and Ba Be and Mg form an adherent oxide coating that allows only slight reaction.
112
Important Reactions of the Alkaline Earth Metals - I
The metals reduce halogens to form ionic halides:
M(s) + X2(-)
MX2 (s)
X = F, Cl, Br, I
Most of the metals reduce hydrogen to form ionic hydrides.
M(s) + H2 (g) MH2 (s)
113
all except Be
Important Reactions of the Alkaline Earth Metals - II Most of the metals reduce nitrogen to form ionic nitrides:
3 M(s) + N2 (g)
M3N2 (s)
all except Be
Except for amphoteric BeO, the oxides are basic:
MO(s) + H2O(l) M(OH)2 (aq) All carbonates undergo thermal decomposition to the oxide: MCO3 (s) MO + CO2 This reaction is used to produce CaO (lime) in huge amounts from naturally occurring limestone, and was the reaction used to generate carbon dioxide to smother the graphite fire in the Chernobyl reactor.
114
General Reactions of Alkaline Earth Metals
115
Diagonal relationship
116
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
117
118
Similarities of Be and Al •
Be and Al have the same electronegativity (Be I.0 and AI 1.5) and their charge/radius ratios are very
•
The standard oxidation potentials of both Be and are of nearly the same order (Be = 1.97V; Al= l.7V)
•
Since the polarizing power of both Be and Al are nearly the same, the covalent character of their compounds also similar.
•
Both Be and Al are rendered ive on treatment with conc. HNO 3.
•
Unlike alkaline earth metals Be does not get readily attacked by dry air. (like Al)
•
Both Be and Al reacts very slowly with dilute mineral acids due to the presence of
oxide layer. •
Both Be and Al react with alkalis liberating H2.
•
Both Be and AI form carbides which on hydrolrolysis liberate methane.
•
Both form nitrides when heated in nitrogen which give ammonia by the reaction with water.
•
Both form oxides which are amphoteric.
119
Halides of both Be and AI contain halogen bridge bonds
•
120
121
Compounds of Na
122
Na2CO3 – Solvay Process
123
124 C a r b o n a t i n g To w e r
125
Sodium Carbonate
126
Caustic Soda – NaOH - Nelson Cell
127
NaOH by Castner Kellner Cell
128
NaOH by Castner Kellner Cell • Castner-Kellner method also known as Mercury Cathode Method. • In this method, the electrolytic cell contains three compartments. (i) Mercury in the outer compartment acts as a cathode while in middle compartment acts as an anode due to induction. (ii) Graphite rods in the outer compartments acts as anode while the iron rods in the middle compartment acts as a cathode.
• Sodium liberated at mercury cathode in the out compartments dissolve in mercury forming sodium amalgam which moves into middle compartment where it react with water at cathode forming NaOH, H2 and Hg. Cl2 gas is liberated at graphite anodes in the outer compartment.
129
Sodium Hydroxide
130
Sodium Hydroxide
131
Sodium Hydroxide
Sodium Beryllate Sodium Aluminate Sodium Stannite Sodium Plumbite Sodium Zincate
132
Sodium Hydroxide
133
Sodium Sulphate - Salt Cake anhydrous Na2SO4 - Glauber’s Salt Na2SO4 .10 H2O •
Preparation
•
The salt cake (anhydrous sodium sulphate) is dissolved in water and the solution is subjected to crystallization.
•
Above 32 °C the anhydrous salt separates.
•
Below 32 °C, the decahydrate salt crystallises out from the aqueous solution.
•
Saturated solution of the decahydrate, on cooling below 12 °C, gives crystals of heptahydrate.
•
Properties
Pb ( NO3 ) 2 Na2 SO4 PbSO4 2 NaNO3 BaCl 2 Na2 SO4 BaSO4 2 NaCl Sr ( NO3 ) 2 Na2 SO4 SrSO4 2 NaNO3
• Uses: It is used in textile industry, medicines as purgative, manufacture of glass plates and sodium salts. 134
Sodium Bicarbonate, Baking Soda, Na2HCO3 • Preparation By ing CO2 through Sodium Carbonate solution but industrially it is manufactured by Solvay's process. • Properties • It is sparingly soluble in Water • Solution is alkaline in nature • Uses
• On heating, it decomposes to give sodium carbonate. • The metal salts which gives basic metal carbonate with sodium carbonate gives normal carbonates.
• Sodium bicarbonate is used, as an antacid in medicine, in dry fire extinguishers, in baking powders and as mild antiseptic for skin infections. 135
Compounds of Alkaline Earth Metals
136
Magneisum Oxide - MgO – Magnesia • Preparation:
1. Calcination of Magnesite (MgCO3) MgCO3 MgO + CO2 2. Heating Mg(NO3)2 or Mg(OH)2
Mg(NO3)2 MgO + 4 NO2 + O2 Mg(OH)2 MgO + H2O
• Properties: • Light infusible white solid • They have high MP 3073K
• Used as refractory material due to the above property 137
Magneisum Oxide - MgO – Magnesia • Chemical properties • Hydrolyses in water to form insoluble Mg(OH)2 2MgO H 2 O 2Mg(OH) 2 • Being basic, reacts with acids to form respective salts MgO 2HCl MgCl 2 H 2 O • Gives Mg on reduction with Carbon at high tempertaures MgO + C Mg + CO
MgO 3C MgC 2 CO • Uses: • Sorel’s cement – used in Dentistry
MgCl2. 5MgO.xH2O
• As an antacid • As an insulator when mixed with asbestos 138
Magnesium Hydroxide Mg(OH)2 • Preparation: 1. By the hydrolysis of MgO 2. By treating MgCl2 with Ca(OH)2 MgCl2 + Ca(OH)2 Mg(OH)2 + CaCl2 • White powdery substance • Sparingly soluble in water
• Used as an Antacid under the name Milk of Magnesia
139
Magnesium Carbonate MgCO3 • Preparation: Hot Magnesium sulfate with sodium bicarbonate
MgSO4 2 NaHCO3 MgCO3 Na2 SO4 H 2O CO2 • Basic magnesium carbonate H 2O
MgSO4 Na2CO3 MgCO3 .Mg (OH ) 2 Na2 SO4 CO2 (Basic Magnesium Carbonate / Magnesia alva)
MgCO3 .Mg (OH ) 2 3CO2 H 2O Mg ( HCO3)2
Mg ( HCO3)2 MgCO3 H 2O CO2
• A solution containing 12% MgCO3 per 100 cc of water containing dissolved CO2 in called Fluid Magnesia 140
Magnesium Sulphate MgSO4.7 H2O Epsom Salt • Preparation • From Magnesite : heating with dil. Sulfuric acid
MgCO3 H 2 SO4 MgSO4 CO2 H 2O • From Dolomite MgCO3.CaCO3 2H 2 SO4 MgSO4 CaSO4 2CO2 2H 2O • From Keiserite (commercial method). Boil with water and cool
MgSO4.H 2O H 2O MgSO4.7 H 2O • Properties • Colourless efflorescent solid C C C MgSO 4 .7H 2 O 30 MgSO 4 .6H 2 O 150 MgSO 4 .H 2 O 200 MgSO 4
141
Magnesium Sulphate MgSO4.7 H2O Epsom Salt • On heating, it decomposes C 4MgSO 4 250 4MgO 2SO 2 2SO 3 O 2
• On heating with Carbon, it gets reduced
2MgSO 4 C 2MgO 2SO 2 CO 2 • It forms double salts with alkali metal sulfates
K 2SO 4 .MgSO 4 .6H 2 O
142
Calcium Oxide / Quick Lime / Burnt Lime • Preparation: Decomposition of Limestone •
C CaCO 3 900 CaO CO 2
• Reacts with water with a hissing noise to form Slaked lime Ca (OH)2 (Rxn is known as Slaking of lime) ΔH = -15 kcal/mol
CaO H 2 O Ca (OH) 2 • Milk of lime : paste of lime in water • Lime water : Clear filtrate • “Limelight” in oxy hydrogen flame
143
Calcium Oxide / Quick Lime / Burnt Lime • CaO gives Calcium Silicate with silica and Calcium Phosphate with P4O10
CaO SiO2 CaSiO3 6CaO 2 P2O5 2Ca3 ( PO4 ) 2 • Forms Calcium Carbide on heating with carbon (2000 deg C) C CaO 3C 2000 CaC 2 CO
• Calcium Carbide + water gives Calcium Cyanamide CaC 2
C N 2 1000
CaCN 2
C
Calcium cyanamide
−
144
Calcium Cyanamide + C = Nitrolim - a Fertiliser
Calcium Oxide / Quick Lime / Burnt Lime • Uses
• It is used as a drying agent. • It is used in the manufacture of bleaching powder. • It is used in the manufacture of calcium. carbide, cement, glass, lime mortar, etc. • It is used in the purification of sugar.
145
Calcium Hydroxide / Slaked Lime / Milk of Lime • Preparation: Slaking of lime • Properties • White amorphous solid • Sparingly soluble in water •
On heating, loses water molecule to form Lime CaO
• Action of CO2
Ca (OH) 2 CO 2 CaCO 3 H 2 O CaCO 3 H 2 O CO 2 Ca (HCO 3 ) 2 Insoluble
So luble
• Similar reaction with SO2 gas is seen when Calcium bisulphite is 146 formed
Calcium Hydroxide / Slaked Lime / Milk of Lime • Reaction with Ammonia Ca (OH) 2 2NH 4 Cl Heat CaCl 2 2NH 3 H 2 O
• Reaction with Chlorine 3Ca(OH)2 + 2Cl2
below 35oC
slaked lime
Ca(OCl)2.Ca(OH)2.CaCl2. 2 H2 O +
H2 O
bleaching powder
•
Bleaching powder is a calcium salt of hypochlorous acid (HOCl)
•
Ca(OCl)2
•
Uses: It is used
1. for absorbing acid gases. 2. in the manufacture of bleaching powder and caustic soda. 3. in the production of lime mortar for construction of buildings, whitewashing buildings 4. in glass making, tanning industry and for purification of sugar 5. for the preparation of NH3 from NH4Cl in Solvay process
147
6. as lime water in laboratories
Gypsum CaSO4.2H2O • Preparation
CaCl 2 H 2SO 4 CaSO 4 2HCl CaCl 2 Na 2SO 4 CaSO 4 2NaCl • Properties • White crystalline solid • Solubility decreases on increase in temperature • Action of heat – Calcium sulphate hemihydrate (plaster of paris) is formed C 2[CaSO 4 .2H 2 O] 120 (CaSO 4 ) 2 .H 2 O 3H 2 O
Plaster of Paris C [(CaSO 4 ) 2 .H 2 O] 200 2CaSO 4 H 2 O
Dead burnt plaster 148
heated 2CaSO 4 Strongly 2CaO 2SO 2 O 2
Calcium Carbonate, CaCO3 • Naturally found as limestone, marble, chalk
• Preparation
Ca (OH) 2 CO 2 CaCO 3 H 2 O
CaCl 2 Na 2 CO 3 CaCO 3 2NaCl • White fluffy powder insoluble in water. But dissolves in water in the presence of carbon-di-oxide to form calcium bicarbonate
CaCO 3 H 2 O CO 2 Ca (HCO 3 ) 2
149
Mortar • It is also known as lime mortar. • It is an intimate mixture of 1 part of slaked lime, 3 parts of sand and water made into paste. • This is used to bind the bricks firmly. • Setting of mortar involves the following steps. • (i) Mortar loses water on of evaporation. • (ii) Carbon dioxide is absorbed from the air converting into calcium carbonate which acts as a binding material. • (iii) Slaked lime reacts with silica forming calcium silicate which gives hardness.
150
Cement • The name portland cement was given to it by Joseph Aspidin (a mason!) because when it is mixed with sand and water it hardens like the lime stone querried at Portland in England. • Composition : CaO 50 to 60 %; SiO2: 20 to 25%; Al2O3 : 5 to 10 %; MgO :2 to 3%; Fe2O3 I to 2% and SO3 1 to 2%. • If lime is excess the cement cracks during setting but if it is less the cement will be weak. • Excess of Al2O3 will make cement quick drying • The raw materials for the manufacture of cement are limestone and alumino silicates (clay, sand and shales). When the powdered raw materials are heated in a rotary kiln, sintered clinker will be obtained. • 151
The setting of cement by mixing with water is due to hydration of the molecules and their rearrangement.
S - Block Metals in Biological Systems
152
Biological functions of Sodium and Potassium ions •
Sodium and potassium are the most common cations in biological fluids.
•
Sodium ion is the major cation of extracellular fluids of animals and in blood plasma, including human beings which is known to activate certain enzymes in the animal body
•
These ions participate in the transmission of nerve signals.
•
They also regulate flow of water across cell membranes and in transport of sugars, amino acids into the cells.
•
Potassium ions are the most abundant cations within cell fluids, where they activate many enzymes that participate in oxidation of glucose to produce adenosine triphosphate (ATP).
•
A typical 70 kg adult contains about 90 g of Na+ ions and 170 g of K+ ions.
•
The daily requirement of sodium and potassium is about 2 g each.
153
Biological functions of Magnesium and Calcium •
Magnesium is an important constituent of chlorophyll.
•
Mg2+ and Ca2+ ions are also responsible for the transmission of electrical impulses along the nerve fibre and the contraction of muscles
•
Calcium ions are essential for the formation of bones and teeth
•
It also plays important roles in maintaining rhythm of heart, clotting of blood, neuromuscular function, interneuronal transmission, cell membrane integrity, etc.
•
The calcium concentration in plasma is regulated at about 100 mg L-1. It is maintained by two hormones, calcitonin and parathyroid hormone.
•
The substance present in bones is continuously solubilized and redeposited to the extent of 400 mg per day in man.
•
All this calcium es through the plasma.
154
The END
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