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Transcript of s BLOCK Elements jee
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s-Block Elements
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The s-block elements
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s-Block Elements
Similarities
Highly reactive metals
Strong reducing agents
Form ionic compounds
Fixed oxidation state
Group I : +1
Group II : +2
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4Variation in Physical Properties of s-block Elements
1. Atomic Radius and Ionic Radius
2. Ionization Enthalpies
3. Hydration Enthalpies
4. Melting Points
5. Electronegativity
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5Atomic 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
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6Atomic Radius and Ionic Radius
Group I
element
Atomic radius
(nm)
Group II
element
Atomic radius
(nm)
Li
Na
K
Rb
Cs
0.152
0.186
0.231
0.244
0.262
Be
Mg
Ca
Sr
Ba
0.112
0.160
0.197
0.215
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
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7On 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
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8Ionisation Enthalpy
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9Ionization Enthalpy
Both atomic radius and ENC down the groupsAtomic radius is more important
IE down the groups
Group I
element1st IE 2nd IE
Group II
element1st IE 2nd IE 3rd IE
Li
Na
K
Rb
Cs
519
494
418
402
376
7 300
4 560
3 070
2 370
2 420
Be
Mg
Ca
Sr
Ba
900
736
590
548
502
1 760
1 450
1 150
1 060
966
14 800
7 740
4 940
4 120
3 390
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Group I
element1st IE 2nd IE
Group II
element1st IE 2nd IE 3rd IE
Li
Na
K
Rb
Cs
519
494
418
402
376
7 300
4 560
3 070
2 370
2 420
Be
Mg
Ca
Sr
Ba
900
736
590
548
502
1 760
1 450
1 150
1 060
966
14 800
7 740
4 940
4 120
3 390
Ionization Enthalpy
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.
For Group II elements, 3rd IE >> 2nd IESimilar reasons can be applied
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Variations in the first and second ionization enthalpies of Group I elements
Ionization Enthalpy
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Variations in the first, second and third ionization enthalpies of Group II elements
Ionization Enthalpy
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Electronegativity
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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
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Group I
element
Electronegativity
value
Group II
element
Electronegativity
value
Li
Na
K
Rb
Cs
1.0
0.9
0.8
0.8
0.7
Be
Mg
Ca
Sr
Ba
1.5
1.2
1.0
1.0
0.9
All have low electronegativity => Electropositive
ELECTRONEGATIVITY
EN down the group
EN : Group II > Group I ( greater ENC)
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Hydration
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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
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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
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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 +
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Group I
ion
Hydration
Enthalpy
(kJ mol1)
Group
II ion
Hydration
enthalpy
(kJ mol1)
Li+
Na+
K+
Rb+
Cs+
519
406
322
301
276
Be 2+
Mg2+
Ca2+
Sr2+
Ba2+
2 450
1 920
1 650
1 480
1 360
Group II > Group I Group II ions have higher charge and small
size higher charge density
stronger ion-dipole interaction
Hydration energies
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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
Melting & Boiling Point
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Group I
element
Melting
Point (C)
Group II
element
Melting
Point (C)
Li
Na
K
Rb
Cs
Fr
180
97.8
63.7
38.9
28.7
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Be
Mg
Ca
Sr
Ba
Ra
1280
650
850
768
714
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
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Group
I
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 -
STRUCTURE
Densities of Li, Na, K
are lesser than that
of water
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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
Anomaly: K is lighter than Na
Li < Na < K < Rb < Cs.
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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.
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Flame Colouration
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Most s-block elements and their compounds give a characteristic flame colour in the flame test
Group I element
Flame colour
Li Crimson
Na Golden yellow
K Lilac
Rb Bluish red
Cs Blue
Flame Colouration
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Most s-block elements and their compounds give a characteristic flame colour in the flame test
Group II element
Flame colour
Be -
Mg -
Ca Brick red
Sr Blood red
Ba Apple green
Beryllium and magnesium atoms are smaller and their electrons being strongly bound to the nucleus are not excited to higher-energy levels.
Flame Colouration
Ca, Sr, Ba
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Mechanism : -
1. In the hotter part of the flame,
2. In the cooler part of the flame,
Na(g) Na(g)*heat
Na(g)* Na(g)cool
[Ne] 3p1 [Ne] 3s1
Ground state
[Ne] 3s1 [Ne] 3p1
+ golden yellow light
Visible region
Flame Colouration
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Mechanism : -
For salts of s-block elements,the metal ions of the salts are first converted to metal atoms
Na+Cl Na(g) + Cl(g)heat
Na(g) Na(g)*heat
Na(g)* Na(g)cool
+ golden yellow light
Na2CO3(s) Na+Cl (more volatile)
Conc. HCl
Flame Colouration
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Complex Formation
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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
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Reasons
1. Absence of low-lying vacant d-orbtals to accept lone pairs from ligands.
For Na+, 1s2, 2s2, 2p6, 3s, 3p, 3dHigh-lying relative to 2p
For Fe2+, 1s2, 2s2, 2p6, 3s2, 3p3, 3d6
Low-lying relative to 3p
Complex formation : Weak Tendency
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Reasons
2. s-block cations (M+, M2+) have relatively low charge densities
less polarizing and less able to accept lone pairs from ligands.
Complex formation : Weak Tendency
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Owing to its high charge density, Be2+ can form complexes
Complex formation :
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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.
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
oEMetallic charater (Reactivity)
Group I > Group II
down the groups
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Reducing Property
Powerful reducing agents
Li > Na < K = Rb > Cs (E0)
Reasons
Heat of sublimation
Ionisation enthalpy
Hydration energy
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Reaction with Hydrogen
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Group I
2M(s) + H2(g) 2MH(s)300C 500C
Alkali metals react with hydrogen to form ionic hydrides M+H-.
The reaction of alkali metals with hydrogen decreases from Li to Cs
Reaction with Hydrogen
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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
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LAH
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
4LiH + AlCl3 LiAlH4 + 3LiClDry ether
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Sodium tetrahydridoborate (sodium borohydride)
NaBH4
Can be used even in aqueous solutions
Na and K hydrides are useful
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Group II
M(s) + H2(g) MH2(s)
600C 700C
Alkaline earth metals react with hydrogen to form ionic hydrides M2+ (H-)2
Reaction with Hydrogen
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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
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Reactions of hydrides
MH(s)
MOH(aq) + H2(g)
MCl(aq) + H2(g)
H (a strong base) tends to react with protonic reagents to release H2
Reactivity down the groups
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Reaction with Air / Oxygen
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Reaction with Air / Oxygen
All alkali metals form more than one type of oxide on burning in air (except lithium)
Group I Elements
All alkaline earth metals react slowly with air to form oxides
On burning in air, they form both oxide and nitride
Group II Elements
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Three types of oxides:
normal oxides
peroxides
superoxides
Reaction with Air / Oxygen : Group 1 Elements
2O
2
1
O2
oxide ion
O22
peroxide ion
2O 2O2
superoxide ion
Abundant supply
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Type of oxide formed depends on
1. supply of oxygen
2. reaction temperature
3. charge density of M+
Reaction with Air / Oxygen : Group 1 Elements
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Lithium
when it is burnt in air, it forms normaloxide only
C180
4Li(s) + O2(g) 2Li2O(s)lithium oxide
Reaction with Air / Oxygen : Group 1 Elements
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Sodium
when it is burnt in an abundantsupply of oxygen
forms both the normal oxide and the peroxide
C180
4Na(s) + O2(g) 2Na2O(s)sodium oxide
C300
2Na2O(s) + O2(g) 2Na2O2(s)sodium peroxideexcess
Reaction with Air / Oxygen : Group 1 Elements
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Potassium, rubidium and caesium
form All three types of oxides when burnt in sufficient supply of oxygen
Reaction with Air / Oxygen : Group 1 Elements
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Group I
elementNormal oxide Peroxide Superoxide
Li
Na
K
Rb
Cs
Li2O
Na2O
K2O
Rb2O
Cs2O
Na2O2
K2O2
Rb2O2
Cs2O2
KO2
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
Reaction with Air / Oxygen : Group 1 Elements
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The electron cloud of the superoxide ion is greatly distorted by the small lithium ion
Reaction with Air / Oxygen : Group 1 Elements
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Group I
elementNormal oxide Peroxide Superoxide
Li
Na
K
Rb
Cs
Li2O
Na2O
K2O
Rb2O
Cs2O
Na2O2
K2O2
Rb2O2
Cs2O2
KO2
RbO2
CsO2
Super oxides are generally bright coloured
They exhibit paramagnetic character due to unpaired electron
Reaction with Air / Oxygen : Group 1 Elements
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KO2 used as oxygen generators and CO2 scrubbers in spacecrafts and submarines
4KO2 + 2H2O 4KOH + 3O2
2KOH + CO2 K2CO3 + H2O
Reaction with Air / Oxygen : Group 1 Elements
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Group II
elementNormal oxide Peroxide
Supero
xide
Be
Mg
Ca
Sr
Ba
BeO
MgO
CaO
SrO
BaO
-
-
All these oxides are basic in nature (except beryllium oxide which is amphoteric)
Reaction with Air / Oxygen : Group 2 Elements
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Solubility
G0 = H0 - TS0
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
LH = 1 / (r+
+ r-) and HydH = (1 / r
+) + (1 / r
-)
Ion size assymmetry results in exothermic dissolution
If both are small, both LH and HydH may be large, but enthalpy of
dissolution may not be very exothermic
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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)
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Two processes are
1. Breakdown of the ionic lattice
2. Hydration
Processes involved in Dissolution and their Energetics
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NaCl(s) Na+(aq) + Cl
-(aq)
Na+(g) + Cl
-(g)
Hsolution
olattice
ohydration
osolution HHH
= (-772 +776) kJ mol1
= +4 kJ mol1
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osolution
osolution
osolution STHG
If , we expect the solids to dissolve in water
0Hosolution
Solubility as becomes more ve (less +ve)osolutionH
Solids (e.g. NaCl) with small +ve valuesare also soluble in water if the dissolution involves an increase in the entropy of the system.
osolutionH
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osolution
osolution
osolution STHG
0Gosolution Spontaneous dissolution
osolutionST is always positive
osolutionHDissolution with slightly positive
can be spontaneous
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Trends and Interpretations
1. The solubility of Group(II) sulphate decreases down the group
On moving down the group, cationic radius(r+)
both and become less -veoLH
ohydrationH
However, less rapidly than oLH
ohydrationH
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Trends and Interpretations
rr
1H
24SO
oL
rr 24SO
constant
olattice
ohydration
osolution HHH
less ve down the group
+ve constantless ve down the group
Solubility down the group
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Trends and Interpretations
rr
1H
24SO
oL
rr 24SO
constant
olattice
ohydration
osolution HHH
more rapidly down the group
less rapidly down the group
less ve down the group
Solubility down the group
(-ve) (+ve)
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Trends and Interpretations
2. The solubility of Group(II) hydroxides increases down the group
On moving down the group, cationic radius(r+)
both and become less -veoLH
ohydrationH
However, more rapidly than oLH
ohydrationH
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Trends and Interpretations
olattice
ohydration
osolution HHH
less rapidly down the group
more rapidly down the group
more ve down the group
Solubility down the group
(-ve)(+ve)
less +ve down the group
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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, CO3
2-),
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
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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
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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
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Thermal Stability
G0 = H0 - TS0
The G0 for the decomposition of a solid becomes negative when
TS0 > H0
H0 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.
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Thermal decomposition reactions
Metal carbonates
M2CO3(s) M2O(s) + CO2heat
MCO3(s) MO(s) + CO2heat
Metal hydroxides
2MOH(s) M2O(s) + H2O(g)heat
M(OH)2(s) MO(s) + H2Oheat
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Relative thermal stability can be measured in two ways
A higher decomposition temperature
a greater thermal stability
C100
BeCO3(s) BeO(s) + CO2(g)
MgCO3(s) MgO(s) + CO2(g) C540
CaCO3(s) CaO(s) + CO2(g) C900
SrCO3(s) SrO(s) + CO2(g) C1290
BaCO3(s) BaO(s) + CO2(g) C1360
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Relative thermal stability can be measured in two waysBy comparing the standard enthalpy changes of thermal decomposition
reactions
A more positive H value a thermally more stable compound
M(OH)2(s) MO(s) + H2O(g) H > 0
Be(OH)2(s) BeO(s) + H2O(g)H = +54 kJ mol1
Mg(OH)2(s) MgO(s) + H2O(g)H = +81 kJ mol1
Ca(OH)2(s) CaO(s) + H2O(g)H = +109 kJ mol1
Sr(OH)2(s) SrO(s) + H2O(g)H = +127 kJ mol1
Ba(OH)2(s) BaO(s) + H2O(g)H = +146 kJ mol1
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Factors affecting thermal stability
1. Polarizing power of cation
2. Polarizability of polyatomic anion
3. Lattice enthalpy of metal oxide produced
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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
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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
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CaCO3(s) CaO(s) + CO2(g)heat
Na2CO3(s) Na2O(s) + CO2(g)
more favourable
less favourable
heat
more stable
less stable
Thermal stability of carbonates : -
Group I > Group II
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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
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MgCO3(s) MgO(s) + CO2(g)heat
more favourable
more stable
BaCO3(s) BaO(s) + CO2(g)heat
less favourable
less stable
more polarized
less polarized
Thermal stability of carbonates
down the groups
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Effect of sizes of the cations on thermal stability of the carbonates and hydroxides of both Groups I and II metals
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Reaction with Water
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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 K
-
Reactions with water or steam
Group I
2M(s) + H2O(l) 2MOH(aq) + H2(g)heat
Group II
M(s) + 2H2O(l) M(OH)2(aq) + H2(g)heat
Mg reacts with steam but not cold water
Be has no reaction with either water or steam
Mg(s) + H2O(g) MgO(s) + H2(g)heat
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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
passing calculated quantities of oxygen gas through the solutions
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Hydroxides
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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
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Group 2: Hydroxides of type M(OH)2
All group 2 metals form hydroxides
Reaction of oxides with water gives hydroxides
Be(OH)2 Mg(OH)2 Ca(OH)2 , Sr(OH)2 , Ba(OH)2
Amphoteric Weakly Basic Strongly Basic
Weaker bases than alkali metal hydroxides
Higher IE, smaller ionic size, higher charge on metal ion.
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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
solubility increases on descending the group.
-
CompoundsSolubility / 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
down the group
CompoundsSolubility / mol per 100 of
water
MgSO4 1800 104
CaSO4 11 104
SrSO4 0.71 104
BaSO4 0.009 104
down the group
-
CompoundsSolubility / 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
CompoundsSolubility / mol per 100 of
water
MgSO4 1800 104
CaSO4 11 104
SrSO4 0.71 104
BaSO4 0.009 104
Size and/or charge of the anion
Polarizability of anion
Covalent character
Solubility in water
In general,
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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 --------------High-------------------------------
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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
state.
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Group II: Carbonates
Thermal Stability Increases down the group
BeCO3 < MgCO3 < CaCO3 < SrCO3 < BaCO3
BeCO3 must be stored under CO2
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96
Reaction with Nitrogen
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97
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 members of N
The binary compounds undergo hydrolysis in water to form
ammmonia, phosphine, asine, stibine etc
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98
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
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99
Halides
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100
Group 1 : Halides
MX
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
2M(s) + Cl2(g) 2MCl(s)heat
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101
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
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102
Halides : Ionic Character
The order of ionic character is
LiX < NaX < KX < RbX < CsX
MF > MCI > MBr > MI
(same metal, different halogen)
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103
Group 2: Halides
MX2
When crystallized from solutions they form hydrated salts
Anhydrous CaCl2, SrCl2 and BaCl2 can be prepared by heating the
hydrated salts.
M(s) + Cl2(g) MCl2(s)heat
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104
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
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105
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
Be atoms, which gives each Be an octet.
Structure of BeCl2 molecules
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106
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
BeF2 > MgF2 > CaF2 < SrF2 < BaF2 BeX2 > MgX2 > CaX2 > SrX2 > BaX2 MF2 < MCl2 < MBr2 < MI2
Saahil Jain
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107
Reactions of chlorides
No significant reactions with water, acids or alkalis
Group I
Group II
Do not undergo significant hydrolysis except BeCl2 and MgCl2
BeCl2(aq) + 2H2O(l) Be(OH)2(aq) + 2HCl(aq)
MgCl2(aq) + H2O(l) Mg(OH)Cl(aq) + HCl(aq)
Basic salt
More favoured in alkaline solutions
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108
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
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109
Group II Sulfates
Solubility in water
BeSO4 > MgSO4 > CaSO4 > SrSO4 > BaSO4
fairly soluble completely insoluble
Thermal stability increases down the group
BeSO4 < MgSO4 < CaSO4 < SrSO4 < BaSO4
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110
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
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111
General Reactions of Alkali Metals
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Important Reactions of the Alkaline Earth Metals - I
The metals reduce O2 to form the oxide:
Barium also forms the peroxide BaO (s).
The Metals of higher atomic weight reduce water to form hydrogen gas:
Be and Mg form an adherent oxide coating that allows only
slight reaction.
2 M(s) + O2 (g) MO(s)
M(s) + 2 H2O(l) M(OH)2 (aq) + H2 (g)M = Ca, Sr and Ba
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113
Important Reactions of the Alkaline Earth Metals - I
The metals reduce halogens to form ionic halides:
Most of the metals reduce hydrogen to form ionic hydrides.
M(s) + X2(-) MX2 (s) X = F, Cl, Br, I
M(s) + H2 (g) MH2 (s) all except Be
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114
Important Reactions of the Alkaline Earth Metals - II
Most of the metals reduce nitrogen to form ionic nitrides:
Except for amphoteric BeO, the oxides are basic:
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.
3 M(s) + N2 (g) M3N2 (s) all except Be
MO(s) + H2O(l) M(OH)2 (aq)
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115
General Reactions of Alkaline Earth Metals
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116
Diagonal relationship
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117
ReactionOther Group I
elementsLithium Magnesium
Combination with O2Peroxides 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
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118
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119
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 passive on treatment with conc. HNO3.
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.
Halides of both Be and AI contain halogen bridge bonds
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120
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121
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122
Compounds of Na
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123
Na2CO3 Solvay Process
-
124
Ca
rb
on
atin
gT
ow
er
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125
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126
Sodium Carbonate
-
127
Caustic Soda NaOH - Nelson Cell
-
128
NaOH by Castner Kellner Cell
-
129
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.
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130
Sodium Hydroxide
-
131
Sodium Hydroxide
-
132
Sodium Hydroxide
Sodium Beryllate
Sodium Aluminate
Sodium Stannite
Sodium Plumbite
Sodium Zincate
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133
Sodium Hydroxide
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134
Sodium Sulphate - Salt Cake anhydrous Na2SO4- Glaubers 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
Uses: It is used in textile industry, medicines as purgative, manufacture
of glass plates and sodium salts.
344223
4422
344223
2)(
2
2)(
NaNOSrSOSONaNOSr
NaClBaSOSONaBaCl
NaNOPbSOSONaNOPb
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135
Sodium Bicarbonate, Baking Soda, Na2HCO3 Preparation
By passing 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.
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136
Compounds of Alkaline Earth Metals
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137
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
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138
Magneisum Oxide - MgO Magnesia
Chemical properties
Hydrolyses in water to form insoluble Mg(OH)2
Being basic, reacts with acids to form respective salts
Gives Mg on reduction with Carbon at high tempertaures
MgO + C Mg + CO
Uses:
Sorels cement used in Dentistry MgCl2. 5MgO.xH2O
As an antacid
As an insulator when mixed with asbestos
22 )OH(Mg2OHMgO2
OHMgClHCl2MgO 22
COMgCC3MgO 2
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139
Magnesium Hydroxide Mg(OH)2
Preparation:
1. By the hydrolysis of MgO
2. By treating MgCl2 with Ca(OH)2
MgCl2 + Ca(OH)2Mg(OH)2 + CaCl2
White powdery substance
Sparingly soluble in water
Used as an Antacid under the name Milk of Magnesia
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140
Magnesium Carbonate MgCO3
Preparation:
Hot Magnesium sulfate with sodium bicarbonate
Basic magnesium carbonate
(Basic Magnesium Carbonate / Magnesia alva)
A solution containing 12% MgCO3 per 100 cc of water containing dissolved
CO2 in called Fluid Magnesia
2242334 2 COOHSONaMgCONaHCOMgSO
24223324 )(.2
COSONaOHMgMgCOCONaMgSOOH
232223 )(3)(. HCOMgOHCOOHMgMgCO
22323)( COOHMgCOHCOMg
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141
Magnesium Sulphate MgSO4.7 H2O Epsom Salt
Preparation
From Magnesite : heating with dil. Sulfuric acid
From Dolomite
From Keiserite (commercial method). Boil with water and cool
Properties
Colourless efflorescent solid
OHCOMgSOSOHMgCO 224423
OHCOCaSOMgSOSOHCaCOMgCO 2244423.3 222
OHMgSOOHOHMgSO 2.4224 7.
4C200
24C150
24C30
24 MgSOOH.MgSOOH6.MgSOOH7.MgSO
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142
Magnesium Sulphate MgSO4.7 H2O Epsom Salt
On heating, it decomposes
On heating with Carbon, it gets reduced
It forms double salts with alkali metal sulfates
232C250
4 OSO2SO2MgO4MgSO4
224 COSO2MgO2CMgSO2
OH6.MgSO.SOK 2442
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143
Calcium Oxide / Quick Lime / Burnt Lime
Preparation: Decomposition of Limestone
Reacts with water with a hissing noise to form Slaked lime Ca (OH)2
(Rxn is known as Slaking of lime) H = -15 kcal/mol
Milk of lime : paste of lime in water
Lime water : Clear filtrate
Limelight in oxy hydrogen flame
22 )OH(CaOHCaO
2C900
3 COCaOCaCO
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144
CaO gives Calcium Silicate with silica and Calcium Phosphate with P4O10
Forms Calcium Carbide on heating with carbon (2000 deg C)
Calcium Carbide + water gives Calcium Cyanamide
Calcium Cyanamide + C = Nitrolim - a Fertiliser
24352
32
)(226 POCaOPCaO
CaSiOSiOCaO
COCaCC3CaO 2C2000
CCaCNNCaCcyanamideCalcium
222C1000
Calcium Oxide / Quick Lime / Burnt Lime
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145
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.
Calcium Oxide / Quick Lime / Burnt Lime
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146
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
Similar reaction with SO2 gas is seen when Calcium bisulphite is
formed
OHCaCOCO)OH(Ca 2322
lelubSolelubInso
)HCO(CaCOOHCaCO 23223
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147
Calcium Hydroxide / Slaked Lime / Milk of Lime
Reaction with Ammonia
Reaction with Chlorine
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
6. as lime water in laboratories
OHNH2CaClClNH2)OH(Ca 232Heat
42
Ca(OCl)2.Ca(OH)
2.CaCl
2. H
2O + H
2O3Ca(OH)2 + 2Cl2
below 35oC
slaked lime bleaching powder
2
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148
Gypsum CaSO4.2H2O Preparation
Properties
White crystalline solid
Solubility decreases on increase in temperature
Action of heat Calcium sulphate hemihydrate (plaster of paris) is formed
HCl2CaSOSOHCaCl 4422
NaCl2CaSOSONaCaCl 4422
OH3OH.)CaSO(]OH2.CaSO[2 2224C120
24
OHCaSO2]OH.)CaSO[( 24C200
224
22heatedStrongly
4 OSO2CaO2CaSO2
Plaster of Paris
Dead burnt plaster
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149
Calcium Carbonate, CaCO3
Naturally found as limestone, marble, chalk
Preparation
White fluffy powder insoluble in water. But dissolves in water in the
presence of carbon-di-oxide to form calcium bicarbonate
OHCaCOCO)OH(Ca 2322
NaCl2CaCOCONaCaCl 3322
23223 )HCO(CaCOOHCaCO
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150
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 account 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.
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151
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.
The setting of cement by mixing with water is due to hydration of the
molecules and their rearrangement.
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152
S - Block Metals in Biological Systems
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153
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.
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154
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 passes through the plasma.
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The END