This document discusses the characteristic properties of s-block elements, which include the alkali metals (Group IA) and alkaline earth metals (Group IIA). Some key points discussed include: - S-block elements have their outermost shell electrons in the s orbital. - Alkali metals react vigorously with water to form alkaline hydroxides and hydrogen gas. Reactivity increases down the group. - They form oxides, peroxides, and superoxides with oxygen. Oxidation states include -2, -1, and -1/2. - Properties such as ionization energy, hydration energy, and metallic character generally decrease or increase moving down a group and across a period,

1. CHARACTERISTIC PROPERTIES OF THE
S-BLOCK ELEMENTS
Dr. Divya Sharma
Assistant Professor

2. PERIODICTABLE OF ELEMENTS

3. S-BLOCK ELEMENTS
• Elements of Groups IA* (the alkali metals) and IIA* (the alkaline earth
metals)
 constitute the s-block elements
 their outermost shell electrons are in the s orbital
• Hydrogen properties are much similar to Alkali earth metals and Halogen family,
so its exception case in Alkali metals.
Na + O₂ → Na₂O
Na + Cl₂ → NaCl
Na + S → Na₂S
H₂ + O₂ → H₂O
H₂ + Cl₂ → HCl
H₂ + S₂ → H₂S

4. Alkali metals Forms oxides and hydroxides
are basic
Alkaline Earth Metals
Form Alkaline oxides
and Hydroxides are
basic
Metal oxides
Group I
Group II
Alkali
Metals
Occur in earth’s curst
Physical Properties

5. Distinguish Between
Group I Metals (Alkali Metal)
vs
Group II Metals (Alkaline Metal)
With respect to their Physical Properties

6. Electronic configuration – Group I
Group I
element
Atomic
number
Electronic
configuration
Electronegativity
value
Oxidation state
in compounds
Li
Na
K
Rb
Cs
Fr
3
11
19
37
55
87
[He] 2s1
[Ne] 3s1
[Ar] 4s1
[Kr] 5s1
[Xe] 6s1
[Rn] 7s1
1.0
0.9
0.8
0.8
0.7
–
+1
+1
+1
+1
+1
–
H Li Na K Rb Cs
2 8 8 18 18 32
Fr

7. Electronic configuration – Group II
Group II
element
Atomic
number
Electronic
configuration
Electronegativity
value
Oxidation state
in compounds
Be
Mg
Ca
Sr
Ba
Ra
4
12
20
38
56
88
[He] 2s2
[Ne] 3s2
[Ar] 4s2
[Kr] 5s2
[Xe] 6s2
[Rn] 7s2
1.5
1.2
1.0
1.0
0.9
–
+2
+2
+2
+2
+2
–
Be Mg
8 8 18 18 32
RaCa Sr Ba

8. Group I
element
Oxide formed
Hydroxide
formed
Flame colour
Li
Na
K
Rb
Cs
Fr
Li2O
Na2O, Na2O2
K2O, K2O2, KO2
Rb2O, Rb2O2, RbO2
Cs2O, Cs2O2, CsO2
–
LiOH
NaOH
KOH
RbOH
CsOH
–
deep red
yellow
lilac
bluish red
blue
–
Oxides and Hydroxides – Group I

9. Oxides and Hydroxides – Group II
Group II
element
Oxide formed
Hydroxide
formed
Flame colour
Be
Mg
Ca
Sr
Ba
Ra
BeO
MgO
CaO
SrO, SrO2
BaO, BaO2
–
Be(OH)2
Mg(OH)2
Ca(OH)2
Sr(OH)2
Ba(OH)2
–
–
–
bluish red
blood-red or crimson
blue
–

10. Less abundant
223
Occurrence Metals
Never found free in nature but
always found in combine state
Reactive Less Reactive
Alkali Metals
Na, K
Li, Rb,
Cs, Fr
21 minutes
Abundant
Radioactive
Fr
half life
Occurrence – Group I

11. 3 2
2 5 2
2 3
2 4 2
3 6
2 4 7 2
2 2
2 2 3 2
i)
ii)
iii)
iv)
v)
vi)
vii)
viii)
ix)
x)
Spodumene, LiAl (SiO )
Petalite, LiAl (Si O )
Common salt or rock salt, NaCl
Sodium carbonate, Na CO
Sodium sulphate or Glauber’s salt, Na SO . 10H O
Cryolite, Na AlF
Borax, Na B O . 10H O
Sylvine, KCl
Carnallite, KCl. MgCl . 6H O
Feldspar, K O.Al O .6SiO
Ores – Group I

12. th
th
Occurrence Metals
Less Reactive
Alkaline earth Metals
Reactive
Ca–5
Mg–6
Sr
Ba
Earth’scrust
Lowerabundance
Be–rare
Ra–Rarest Radioactive
Occurrence – Group II

13. 3
2 2
4 2
4
4 2
2 4 2
3 4 2
i) Magnesite, MgCO
3
ii) Dolomite, MgCO .CaCO
3
iii)
iv)
v)
vi)
vii)
Carnallite, KCl.MgCl .6H O
Epsom salt (Epsomite), MgSO .7H O
Limestone, marble, chalk or calcite, CaCO
Anhydrite, CaSO
Gypsum, CaSO .2H O
3
viii) Fluorspar, CaF
2
ix)
x)
Hydroxyapatite, 3 Ca (PO ) .Ca(OH)
Phosphorite, Ca (PO )
2
Ores – Group II

14. Li
Na
K
Rb
ATOMIC
RADII
Ca
Sr
Mg
–
–
–
–
–
–
– – –
– – –
– –
–
Group 1 Group 2
– e
e
e
e
e
e
e
e
e
e
Across the period
atomic size
decreases
Down the group
atomic size
increases
e
e e e
e
e
Atomic radius
Be
Mg
Ca
Sr

15. IONIC RADII
Li Li+
Be
–
–
–
–
–
–
–
–
–
–
–
Down the group
e e
e
>
e
e
152 pm 76 pm
e e
e
e >
e
e
Ionic size
increases
112 pm 31 pm
Be Be+2
Ionic radius

16. Atomic radius & Ionic radius
Group I element Atomic radius (nm) Ionic radius (nm)
Li
Na
K
Rb
Cs
0.152
0.186
0.231
0.244
0.262
0.060
0.095
0.133
0.148
0.169
Group II element Atomic radius (nm) Ionic radius (nm)
Be
Mg
Ca
Sr
Ba
0.112
0.160
0.197
0.215
0.217
0.031
0.065
0.099
0.113
0.135

17. The ionization energy (IE) is qualitatively defined as the amount of
energy required to remove the most loosely bound electron, the valence
electron, of an isolated gaseous atom to form a cation.
Ionization Energy
1 Ionization energy of group I elements are lower
than 1 Ionization energy of group II elements. Due
to size decreases required more energy.
st
st
2 Ionization enthalpy of group I elements are
greater than 2 Ionization enthalpy of group II
elements.
As after losing the 1 electron from outer most shell,
group I become more stable due to noble gas
configuration than group II element.
I.E.1 < I.E.1
Na < Mg
I.E.2 < I.E.2
Na < Mg
Na = 1s2, 2s2, 2p6
Mg = 1s2, 2s2, 2p6, 3s1
nd
nd
+ +
+
+

18. Li
Na
K
Be
Mg
Ca
– –
–
–
–
–
–
– –
–
–
–
–
e
Group 1 Group 2
e
e
e
e
e
e
e
e
e
e
Ionization
Energy
Down the group
IE decreases
e
Stable configuration
Caesium used in photoelectric cell
e
Across the period
IE increases

19. Group I element
First ionization
enthalpy (kJ mol–1)
Second ionization
enthalpy (kJ mol–1)
Li
Na
K
Rb
Cs
Fr
519
494
418
402
376
381
7 300
4 560
3 070
2 370
2 420
–
Group II
element
First ionization
enthalpy (kJ mol–1)
Second ionization
enthalpy (kJ mol–1)
Third ionization
enthalpy (kJ mol–1)
Be
Mg
Ca
Sr
Ba
Ra
900
736
590
548
502
510
1 760
1 450
1 150
1 060
966
979
14 800
7 740
4 940
4 120
3 390
–

20. +
With Na and H O Interaction
–
+
2
δΗ
δΗ δΗ
+ +
δΗ
+ +
δ H
+ O δ
–
O δ
O δ
–
δΗ
+
Hydration reaction is a chemical
reaction in which a substance
combines with water. This Process is
known as hydration.
Hydration enthalpy (Hhyd) is the
amount of energy released when one
mole of aqueous ions is formed from
its gaseous ions.
O
δ
–
Na+
δΗ
+
δΗ
+
O δ
H+
δ
–
O δ
–
δΗ
+
H
δ
δΗ
+
Hydration Energy
M+(g) + aq  M+(aq)
H = Hhyd
 always has a negative value

21. +
With Na and H O Interaction
2
This is hydrated sodium ion.
Hydration energy
The energy involved in this process is called
H O
2
H O
2
Na
H O2
+
H O2
Hydration energy.
1
Hydration energy ∝
Size of cation
H O H O2 2
Hydration Energy
Size of Cation Hydration
energy
Smallest Li+ Maximum
Na+
K+
Rb+
Cs+
Biggest Fr+ Minimum
DecreasesIncreases

22. Group I ion
Hydration enthalpy
(kJ mol–1)
Group II ion
Hydration enthalpy
(kJ mol–1)
Li+
Na+
K+
Rb+
Cs+
Fr+
–519
–406
–322
–301
–276
–
Be 2+
Mg2+
Ca2+
Sr2+
Ba2+
Ra2+
–2,450
–1,920
–1,650
–1,480
–1,360
–
Hydration Energy

23. Alkali metal Alkaline earth metal
7]
Oxidation
state + 1 + 2 Li+ ion is the strongest
reducing agent
8]
Reducing
property
EXCELLENT
reducing agent
WEAKER reducing agent
compared to alkali metal
Na < K < Rb < Cs < Li
The reducing character
of any metal is best
measured in terms of
its electrode potential
1) Heat of vapourization
2) Ionization enthalpy
3) Heat of hydration
Oxidatio
n State
Reducin
g Agent
Oxidation and Reducing agent

24. Alkali metal Alkaline earth metal
1 +
M
9]
MOST metallic
due to alkali metals are more
electropositive in nature
LESS metallic
Metals have a great tendency to
loose. These metals are highly
electropositive in nature.
ns
electron ions
Group 2 elements are less
electropositive than Group 1
elements
Down the group Electropositive /
Metallic Character INCREASES
Metallic
Characte
r
Electropositi
ve Character
10]
1Metallic
Character ∝
Across a period Electropositive /
Metallic Character DECREASES
Metallic and Electropositive Character
Ionization energy

25. Alkali metal Alkaline earth metal
Down the group
melting point
DECREASES
Group 2 elements have higher
M.P and B.P than Group 1
elements
M – N.M → Ionic bond
Liquid
Melting
Point
&
Boiling Point
11]
Melting & Boiling Point
Regular pattern of M.P. & B.P.
Li > Na > K > Rb > Cs
1
M.P. & B.P. ∝
Atomic Size
Irregular pattern
M.P. = Be > Ca > Sr > Ba > Mg
B.P. = Be > Ba > Ca > Sr > Mg
N.M – N.M → Covalent bond
M – M → Metallic bond

26. Alkali metal Alkaline earth metal
Low
density
Alkaline earth
metals are DENSER
than alkali metals
 Down the group atomic mass increases,
which more than compensates the bigger
atomic size therefore Density Increases.
 K, however, is lighter than Na which is due
to an abnormal increase in its atomic size
because of presence of penultimate d-orbitals.
 Ca is less than Mg because of presence of
penultimate d-orbitals.
Lightest metal
Density12]
Density

27. Alkali metal
Crimson red
Golden yellow
Pale violet
Purple
Sky blue
(violet)
Fr being radio active does
not give flame colouration
Li Na K Rb Cs
Flame
coloration
Flame coloration
 Both groups show flame coloration.
 The electron present in an atom absorb energy from flame and there electrons
are excited to higher energy level. When these electron fall back to ground state
emit radiation of different wavelength imparting color to the flame.

28. Brick
red
Crimson
red
Apple
green
Crimson
Alkaline earth metal
Ca Sr Ba Ra
 Be and Mg because of their high
ionization energy, however, do
not impart any characteristic
color to the Bunsen flame.
 These property generally shows
only in s-block elements.
d-block elements – do not this
flame coloration property.
Flame
coloration

29. Chemical Properties and General
Characteristics of
Group I Metals
(Alkali Metal)

30. 1) Reaction with water (moisture)
2 M + 2 H O
2
2 MOH + H
2
e.g :
2 Li + 2 H O
2
2 LiOH + H
2
2 Na + 2 H O
2
2 NaOH + H
2
Down the group
Reactivity
INCREASES
Most reactive :
Least reactive :
Cs
Li
LiOH is a weak base,but all other
hydroxides are highly basic
Chemical Properties – Group I
Chemical Property (Alkali metal)

31. M O
1 2
KO + H O KOH + O
KOH + CO KHCO
2 1
2
–2
Na O O2 2
– +...
2
2
2(a). Reaction with air (oxygen)
Metal + Oxygen Metal oxide
M + O
2
Li Li + O
2
M O M
Li O
2
M O
Metal + Oxygen Metal peroxide
M + O
2
Na Na + O
2
M O O M
O O
(M O )2 2
2
2 3
2
Metal + Oxygen Metal superoxide
M + O (O O ) M
(MO )
2
K , Rb , Cs K + O K O
2
(O...O )–
2
Chemical Properties – Group I
O
-2
–1
O2
Oxidation State
Oxide = -2
Peroxide = -1
Superoxide = -
1/2

32. –2
–2 +
+
–1 +
+
O
O
2
Oxide
ion
Peroxide
ion
Li
Na
+
Small cation will be Stabilized
by small anion
K
Only lithium forms oxide
i.e. monoxide
O
2
Superoxide
ion
Rb
Cs
Chemical Properties – Group I

33. M N
2(b). Reaction with air (Nitrogen)
Metal + Nitrogen ⇒ Nitride
M
6M + N
2
Li
M N M
2M N
3
1 3
6Li + N
2
2 Li N
3
Only Li forms nitride
Chemical Properties – Group I

34. M H
1 1
3) Reaction With Hydrogen
2M + H
2
2 M H
Down the group
Reactivity
INCREASES
 Li 1073 K Other Alkali Metals 673 K
 LiH is the MOST stablest of all hydrides
4) Reaction With Halogen
LiF is insoluble in water
2M + X
2
2 M X
Reactivity : M ⇒ Li > Na > K > Rb > Cs (Depends on Size)
X ⇒ F > Cl > Br > I
Chemical Properties – Group I

35. + –
5) Solubility in Liquid Ammonia
All the alkali metals dissolve in liquid ammonia giving Deep-blue solutions
when dilute, due to the presence of ammoniated (solvated) electrons in the
solution
M +(x + y) NH
3
M (NH ) + e (NH )
3 x 3 y
These electrons are excited to higher energy levels and the ABSORPTION of
photons occurs in the RED region of the spectrum
Thus the solution appears blue
Chemical Properties – Group I

36. Solutions of alkali metals (Li, Na and K) in liquid ammonia :
1. Conducting due to the presence of ammoniated electrons mainly.
▪ However, on heating, the conductivity increases (On raising the
temperature of the solution)
▪ Conductivity decreases like those of metals
2. Strongly reducing due to the presence of ammoniated electrons and
are widely used as reducing agents in organic chemistry under the
name BIRCH REDUCTION
Chemical Properties – Group I

37. General characteristics of the compounds of the alkali metals
▪ All the common compounds of the alkali metals are generally
IONIC in nature
Oxides and Hydroxides
+ -
+ –
+ –
❑ Under appropriate conditions pure compounds M2O, M2O2 and MO
may be prepared.
2
M2O + H2O → 2M + 2OH
M2O2 + 2H2O → 2M + 2OH + H2O2
2MO2 + 2H2O → 2M + 2OH + H2O2 + O2
These oxides are easily hydrolyzed by water to form the hydroxides
General Characteristics – Group I

38. The superoxides are paramagnetic
–
The oxides and the peroxides are colorless when pure, but the superoxide's
are yellow or orange in colour
Example : KO
2
The superoxide O2 is paramagnetic because of One Unpaired
Electron in π*2p molecular orbital
The alkali metal hydroxides are the STRONGEST of all bases and dissolve freely in
water with evolution of much heat on account of intense Hydration
General Characteristics – Group I

39. Halides
The low solubility of LiF in water is due to its high lattice
enthalpy
the low solubility of CsI is due to smaller hydration enthalpy
of its two ions.
The alkali metal halides, MX, (X=F, Cl, Br, I) are all high
melting, colorless CRYSTALLINE solids.
• All these halides are soluble in water.
• The melting and boiling points always follow the trend:
FLUORIDE > CHLORIDE > BROMIDE > IODIDE
whereas
General Characteristics – Group I

40. Salts of Oxo-Acids
What are Oxo-Acids?
Oxo-acids are those in which the acidic proton is on a hydroxyl
group with an oxo group attached to the same atom.
Example : Carbonic acid, H CO (OC(OH)
2 3 2
Sulphuric acid, H SO (O S(OH) )2 4 2 2
The alkali metals form Salts with All the oxo-acids.
General Characteristics – Group I

41. Carbonates
M CO MHCO
They are generally Soluble in water and Thermally stable.
Their and
Hydrogen carbonates
2 3
Are highly stable to heat.
3
Lithium carbonate
is Not So Stable to heat
Lithium being very small in size polarises a large CO3 ion leading to the
formation of more stable Li O and CO .
Hydrogen carbonate
Does Not Exist as a solid
2 2
2
-
Li + CO3 -- Li2O + CO2
2-+
General Characteristics – Group I

42. radius
SMALL size
charge
Highest POLARIZING power i.e. ratio
High IONIZATION enthalpy
Low ELECTROPOSITIVE character compared to other metal
ANOMALOUS PROPERTIES OF LITHIUM

43. LithiumAlkali Metals
Water 2 M + 2 H O
2
2 MOH + H 2 2 Li + 2 H2O 2 LiOH + H2
Strong base Weak base
Air
Na + O Na O
2 2 2
Li + O
2
Li O
2
Hydrogen
K + O
2
2M + H
2
KO
2
2 MH
6Li + N
2
LiH
2 Li N
3
Unstable Stable
Halogen 2M + X
2
2 MX
Soluble
LiF
Insoluble
Metal + Oxygen Metal
Oxide
Oxide
Superoxi
de
Lithium Oxide
Lithium
Nitride
ANOMALOUS PROPERTIES OF LITHIUM

44. ANOMALOUS PROPERTIES :
LiOH
Li
Weak base
least reactive
Other alkali metal from strong hydroxides.
Li + Air
Strong reducing agent
Li O (Lithium monoxide)
2
Li N (Lithium nitride)
3
Other alkali metals do
not form monoxide
and nitrides
LiH most stable Other alkali metal hydrides are not much stable
Lithium carbonate
Lithium fluoride
Lithium phosphate
(Li CO )
2 3
(LiF)
(Li PO )
3 4
Sparingly soluble
2
(In Periodic
table)
(Less soluble in water)
ANOMALOUS PROPERTIES OF LITHIUM
But corresponding salts of other alkali
metals are soluble in H O

45. 5) Li
M.P and B.P
6) LiCl
Hard metal
high
Deliquescent
To dissolve and become liquid by
absorbing moisture from airDue to high hydration enthalpy
Crystallize LiCl.2H O
2
s
Other alkali metals chlorides
7)
do not form hydrates
Δ
4 Li NO 2Li O + 4 NO + O
3 2 2 2
(Lithium nitrate)
2 Na NO
3
(Sodium nitrate)
Δ
(Lithium monoxide)
2 NaNO + O
2 2
(Sodium nitrite)
Due to hard metallic crystalline
structure – high lattice enthalpy
ANOMALOUS PROPERTIES OF LITHIUM

46. Li2 NH + H
2 2
8) Li CO
2 3
(Lithium
Δ Li O + CO
2 2
Carbonate)
Other alkali metal carbonates do not evolve carbon dioxide
9) Lithium hydrogen carbonate (LiHCO )
3
Not obtained in solid form
Other alkali metal hydrogen carbonates Solid form
10) Li + NH
3
Other alkali metal form amides (MNH )
2
M = Na, K, Rb, Cs
2 Na + 2 NH
3
Fe (NO )
3 3
Ferric nitrate
2 NaNH + H
Lithium Imide
Soda amide
ANOMALOUS PROPERTIES OF LITHIUM

47. Diagonal Relationship
A diagonal relationship is said to exist between
certain pairs of diagonally adjacent elements in
the SECOND and THIRD Periods of the
periodic table.
DIAGONAL RELATIONSHIP

48. Atomic Radii :
Li = 152 pm
Mg = 160 pm
Ionic Radii :
Li = 76 pm
Mg = 72 pm
Li 1.7
Mg = 3.9
Electronegativities::
Li = 1.00
Mg 1.20
Polarizing power :
+
Diagonal Relationship Similarity b/w Li and Mg
REASON :
Similarity in ionic sizes
charge
Similarity in
radius
ratio
2+
+
++
DIAGONAL RELATIONSHIP

49. 2
3 3 2
2 2 2
2 3
3
Points of Similarity between Li and Mg
1) LiOH is a WEAK base 1) Mg (OH) WEAK base
2) Li → Li N
3) Lithium is HARD
2) Mg N
3) Even Magnesium is HARD
4) LiCl deliquesent 4) Even MgCl
2
LiCl. 2H O MgCl . 8H O
5) Li CO → Decomposes 5) MgCO also Decomposes
6) Lithium hydrogen carbonate 6) Magnesium hydrogen carbonate
cannot be obtained in solid form cannot be obtained in solid form
DIAGONAL RELATIONSHIP

50. Chemical Properties and General
Characteristics of
Group II Metals
(Alkaline Earth Metal)

51. CHEMICAL PROPERTY (ALKALINE EARTH METAL)
1. Reaction with air (Oxygen)
(M O)
M2 O2+ Oxygen
M + O2
Be + O
2
Metal Metal oxide
Ba, Sr, Ra
M + O2 MO 2
O O
O2
–2
Be, Mg and Ca when heated with oxygen form monoxides while Sr, Ba
and Ra form peroxide
M O
Be O
Mg O
M O
2 2
Chemical Properties – Group II
Mg + O2

52. 2. REACTION WITH NITROGEN
+3M + N2
M3N2
Be , Mg
N
M
M
N
M
M N
2 3
2
3 Be + N
Mg + N2
Be 3N 2
Mg N3 2
The ease of formation nitrides increases from Be to Ba.
These nitrides react with water to evolve NH3
Mg3 N2 + 6H2O → 3 Mg (OH)2 + 2 NH3
M
3
N
2
Chemical Properties – Group II

53. OH
OH
Down the group Reactivity INCREASES
Be Does not react with boiling water
Reacts with boiling waterMg
Ca, Sr & Ba Reacts vigorously even with cold water
3. REACTION WITH MOISTURE ( WATER)
2
2M + 2H O M (OH) + H
2 2
M
Chemical Properties – Group II

54. 4. REACTION WITH HYDROGEN
M H2
Mg H2
CaH2
2BeCl2 + LiAlH4 → BeH2 + LiCl +AlCl
3
BeH can, however be prepared by2
reducing BeCl2 with LiAlH4
M + H2
Mg + H2
Ca +
H2 Hydrolith
All the hydrides react with water to evolve H2 and thus behave as
strong REDUCING AGENTS.
MH2 + 2 H2O → M (OH)2 + 2 H2
Chemical Properties – Group II

55. SOLUTIONS IN LIQUID AMMONIA
Like alkali metals, the alkaline earth
metals dissolve in liquid ammonia to give deep
blue black solutions forming ammoniated ions.
M + (x + y) NH3 → [M(NH3)x ] 2+ + 2 [e(NH3)y ]-
3 6
From these solutions, the ammoniates, [M(NH ) ] 2+ can be recovered but
concentrated solution are bronze coloured due to the formation of metal clusters
These solution are good conductors of electricity
Chemical Properties – Group II

56. GENERAL CHARACTERISTICS OF COMPOUNDS OF THE ALKALINE-EARTH METALS
Oxides and Hydroxides
□ The alkaline earth metals form compounds which are predominantly
ionic but less ionic than the corresponding compounds of alkali metals
▪ The alkaline earth metals burn in oxygen to form the monoxide,
MO.
BeO is amphoteric while oxides of other elements are basic in nature.
□ All these oxides except BeO are basic in nature and react
with water to form sparingly soluble hydroxides.
MO + H2O → M(OH)2
General Characteristics – Group II

57. □ The Solubility, Thermal Stability and the Basic Character of these hydroxides
Increases with increasing atomic number from Mg(OH)2 to Ba(OH)2.
□ The alkaline earth metal hydroxides are, however, less basic and
less stable than alkali metal hydroxides.
General Characteristics – Group II

58. Halides
Except for beryllium halides, all other halides of alkaline
earth metals are IONIC in nature
Beryllium chloride has a chain structure In the solid state as shown below:
General Characteristics – Group II

59. IN THE VAPOUR PHASE BECl2 TENDS TO FORM A
CHLORO-BRIDGED DIMER
which dissociates into the linear monomer at high
temperatures of the order of 1200 K.
General Characteristics – Group II

60. SALTS OF OXOACIDS
The alkaline earth metals also form salts of oxoacids.Carbonates
Carbonates of alkaline earth metals are insoluble in water and
can be precipitated by addition of a sodium or ammonium
carbonate solution to a solution of a soluble salt of these metals
The solubility of carbonates in water DECREASES as the
atomic number of the metal ion increases
▪ All the carbonates decompose on heating to give carbon dioxide
▪ Beryllium carbonate is unstable and can be kept only in
the atmosphere of CO2
General Characteristics – Group II

61. SULPHATES
▪The sulphates of the alkaline earth metals are all white solids and
stable to heat.
• BeSO4 and MgSO4 are readily soluble in water.
• The SOLUBILITY decreases from CaSO4 to BaSO4
▪ The greater hydration enthalpies of Be2+ and Mg2+ ions overcome the lattice
enthalpy factor and therefore their sulphates are soluble in water.
General Characteristics – Group II

62. NITRATES
The nitrates are made by dissolution of the carbonates in dilute
nitric acid
▪ Magnesium nitrate crystallises with six molecules of water, whereas
barium nitrate crystallises as the anhydrous salt.
This again shows a decreasing tendency to form hydrates with
INCREASING SIZE and DECREASING HYDRATION
ENTHALPY.
All of them decompose on heating to give the oxide like
Lithium nitrate.
General Characteristics – Group II

63. ANOMALOUS BEHAVIOUR
Higher
electronegativity

64. i) Be is hard metal other alkaline metals are soft
ii) Be Least metallic, forms covalent compounds
Easily hydrolysed.
iii) M.P, B.P, I.E
iv)
BeO and Be(OH)2
Other oxides and hydroxides
v) Be Does not liberate H2 from
Basic in nature
HCl, H2SO4
other alkaline metals liberate H2
vi) Be Does not show coordination number more than
4 d orbitals are absent.
Highest of all the alkaline earth metals
Amphoteric
Anomalous Properties of Beryllium

65. Some Examples of Group I Metals
with their properties and uses

66. SODIUM CARBONATE1)
a) Chemical name
Common name
Molecular formula
= Sodium carbonate
= Washing soda
= Na2CO3 . 10H2O
b) Preparation :
Process
Solvay process
(Ammonia soda process)
Raw materials
=
= NaCl, CaCO3, NH3, water etc.
Sodium Carbonate

67. Sodium bicarbonate
+ H2O + CO22 NaHCO3 Na2CO3
NaHCO3 + NH4Cl
NH3 + H2O + CO2
NaCl + NH4HCO3
NH4HCO3
Sodium Carbonate

68. □ This process is strictly applicable for NaHCO3 and not
KHCO3.
□ NaHCO3 is precipitated due to common ion (Na+)
▪ NaHCO3 is sparingly soluble in H2O hence can be
precipitated
▪ KHCO3 is highly soluble hence can not be precipitated
Sodium Carbonate

69. Na2CO3 . 10H2O
Na2CO3 . H2O above 373K
C) PROPERTIES
1) It’s a white, crystalline substance.
It is readily SOLUBLE in water.
Sodium carbonate on heating ,
2)
3)
Na2CO3 . 10H2O below 373K Na2CO3 . H2O +9H2O
2 3 2
Na CO + H O
4) +
(carbonate ion)
Na2CO3 2Na + CO3
Soda ash
2–
Na2CO3 + 2H2O H2CO3 + 2NaOH
(weak acid) (strong base)
2
H O
Sodium Carbonate

70. d) Uses
1) Soap, detergents, paper, borax and caustic soda.
2) Water softening, laundering and cleaning.
3) Paints and Textile industries.
4) Qualitative and Quantitative analysis.
Sodium Carbonate

71. SODIUM
BICARBONATE
2)
= Sodium bicarbonate
= Baking soda
= NaHCO3
a) Chemical name
Common name
Molecular formula
b) Preparation :
Process
= Solvay process
(Ammonia soda process)
Na2CO3 + H2O + CO2 2 NaHCO3
Sodium Bicarbonate

72. NaHCO3 + HCl NaCl + H2O + CO2
1) It’s a white, crystalline substance.
2) 2NaHCO3 Na2CO3 + H2O + CO2
3) Reaction with acid
4) It’s aqueous solution is ALKALINE due to hydrolysis
NaHCO3 + H2O H2CO3 + NaOH
(weak acid) (strongbase)
C) Properties
Sodium Bicarbonate

73. Fire extinguisher
Baking powder
Antacid
Qualitative and Quantitative analysis.
1)
2)
3)
4)
D) Uses
Baking powder = 30% NaHCO3 + 40% starch + 10 %
calcium dihydrogen phosphate + sodium Al
sulphate
Sodium Bicarbonate

74. SODIUM CHLORIDE
a) Chemical name
b) Common name
c) Molecular formula
d) Preparation Process:
3)
= Sodium chloride
= Table salt or common salt
= NaCl
=
Sodium Chloride
Evaporation by sea water

75. 1) It’s a white, crystalline substance.
2) Melting point 8010C
3) SOLUBLE in water (Solubility of 36.0 g in 100g ofwater at 273 K)
and SLIGHTLY SOLUBLE in alcohol
C) Properties
Sodium Chloride

76. 1) Essential CONSTITUENT of our diet
2) Also used as a FOOD PRESERVATIVE
3) FREEZING MIXTURE to lower down temperature of ice
4) Preparation of NaOH, Na2O2 and Na2CO3.
5) Salt + little Na2CO3 + 5 to 10% Na2SO4 + Sugar = KALA
NAMAK Digestion
D) Uses
Sodium Chloride

77. SODIUM HYDROXIDE4)
Chemical name
Common name
Molecular formulae
Preparation :
Process
= Sodium hydroxide
= Caustic soda
= NaOH
= Electrolysis of NaCl (brine)
Castner – Kellner Cell
Mercury – Cathode Cell
Sodium Hydroxide

78. Brine solution
Anode
Cathode
Electrolyte
= Graphite rods
= Mercury
=
At anode :
Cl– Cl + Cl Cl2
Na / Hg
Cl +e–
At cathode :
Na+ + e– Na Na + Hg
(Sodium amalgam)
(Na+Cl–)
The amalgam is treated with water
2Na / Hg + 2H2O 2NaOH + H2+ 2Hg
Sodium Hydroxide

79. c) Properties
1) It’s a white, translucent, deliquescent solid.
2) Melting point 591 K.
3) Highly SOLUBLE in water.
4) The solution of NaOH at the surface
reacts with CO2 in the atmosphere
to form Na2CO3.
2NaOH + CO2 Na2CO3 + H2O
Sodium Hydroxide

80. 3)
4)
1) Purification of bauxite
2) Manufacture of soap, paper, artificial silk and a
number of chemicals.
3) Petroleum refining
4) In the TEXTILE industries for Mercerizing cotton
fabrics.
5) Laboratory REAGENT
6) In preparation pure fats and oils
d) Uses
Sodium Hydroxide

81. Some Examples of Group II Metals
with their properties and uses

82. CALCIUM CARBONATE1)
a) Chemical name
Common name
Molecular formula
= Calcium Carbonate
= Lime stone
= CaCO3
b) Preparation :
Process
Chemical reaction=
Calcium Carbonate

83. Lab method
1) By passing carbon dioxide through slaked lime
Ca(OH)2 + CO2 CaCO3+H2O
2) By adding sodium carbonate to calcium chloride
CaCl2 +Na2CO3 CaCO3 +2NaCl
CaCO3 +H2O+ CO2 Ca(HCO3)2
(insoluble)
(soluble)
Calcium bicarbonate
Calcium Carbonate
Controlled addition of CO2 is essential

84. c) Properties
1) White fluffy powder.
2) Insoluble in water.
3) On heating
CaCO3 CaO + CO2
4) It reacts with dilute acid to
CaCO3 + 2HCl CaCl2 + H2O + CO2
CaCO3 + H2SO4 CaSO4 + H2O + CO2
Calcium Carbonate

85. d)Uses
1) For building materials in the
form of (quick lime, cement,
marble, distemper)
2) Used as a flux in the extraction of
metals
SiO2 + CaCO3 CaSiO3 + CO2
3) Specially precipitated CaCO3 is
extensivelyused
Calcium Carbonate

86. CALCIUM OXIDE2)
a) Chemical name
Common name
Molecular formula
= Calcium Oxide
= Quick lime
= CaO
b) Preparation :
Process
Heating limestone=
Calcium Oxide

87. PREPARATION :
Commercially by heating CaCO3 in a reverberatory kiln at 1070 – 1270 K
CaCO3 CaO + CO2
heat
Crushed
limestone
Good
draught
of air
CO2
Fuel
Rotary kiln
lime
Calcium Oxide

88. c) Properties :
1) White AMORPHOUS solid.
2) Melting point of 2270 K.
3) On exposure to atmosphere it reacts with CO2
CaO + CO2 CaCO3
4) Being BASIC OXIDE combines with acidic oxides at high temperature.
CaO + SiO2 CaSiO3
6CaO + P4O10 2Ca3(PO4)2
5) On heating with ammonium salt it gives ammonia.
CaO + 2NH4Cl CaCl2 + 2NH3 + H2OΔ
Calcium Oxide

89. 1. Manufacturing CEMENT.
2. Cheapest form of ALKALI.
3. Manufacturing of sodium
carbonate from caustic soda.
4. Purification of sugar.
5. Manufacturing of DYE
STUFFS.
D)USES
Calcium Oxide

90. + CLAY
+
GYPSUM
CaSO4 . 2H2O
Calcium Oxide

91. Plaster of Paris
Calcium Sulphate Hemihydrate
Calcium Oxide

92. Plaster of Paris
CaSO
4 2 22 2
. 1 H O + 1 1 H O Setting
HardeningCaSO4 . 2H2O
Gypsum
(orthorhombic)
CaSO4 . 2H2O
Gypsum
(monoclinic)
PLASTER OF PARIS
Calcium Sulphate Hemihydrate
 The setting of Plaster of Paris is believed to be due to
rehydration and its reconversion into gypsum.
Calcium sulphate Hemihydrate

93. Na = 90 g
K = 5 g
Mg = 25 g
Ca = 1200 g
containing only 5g of iron and 0.06 g of copper
The metal required in major
proportions are s-block elements
23 ESSENTIAL ELEMENTS
ARE REQUIRED OUT OF
THAT 15 ARE METALS
Calcium sulphate Hemihydrate

94. HUMAN CELL
BLOOD
PLASMA
High concentration of Na+
and low concentration of K+
High concentration of K+ and
low concentration of Na+
Na = 10 mg L–1
K = 105 mg L–1
Na = 143 mg L–1
K = 5 mg L–1
Inside the cell and tissue
In fluid, bathing cells
and blood plasma
Calcium sulphate Hemihydrate

95. Sharma’s Classes - Dr. Divya Sharma