Chapter 10 s & p Block Elements part 1 by TEACHING CARE Online coaching and tuition classes

Chapter 10 s & p Block Elements part 1 by TEACHING CARE Online coaching and tuition classes



(1)  Position of hydrogen in the periodic table

Hydrogen is the first element in the periodic table. Hydrogen is placed in no specific group due to its


property of giving electron (When

H is formed) and also losing electron (When

H + is formed).


  • Hydrogen is placed in group I (Alkali metals) as,


  • It has one electron in its (Outer) Shell-

1s1 like other alkali metals which have (inert gas)



  • It forms monovalent
  • It valency is also

H + ion like

Li + , Na +


  • Its oxide (H 2 O) is stable as Li2 O, Na2 O .


  • It is a good reducing agent (In atomic as well as molecular state) like
    • Hydrogen also resembles halogens (Group VIII A) as,

Na, Li…


  • It is also diatomic (H 2 ) like F2 , Cl2 …


  • It also forms anion

H like

F , Cl … by gain of one electron.


  • H has stable inert gas (He) configuration as CH 4 , C2 H6 like halogens CCl4 , SF2 Cl2


  • H is one electron short of duplet (Stable configuration) like deficient than octet, F – 2s 2 2p5 ; Cl – 3s 2 3p5 .

F, Cl,… which are also one electron


  • (IE) of H(1312 kJ mol -1 ) is of the same order as that of
    • (IE) of H is very high in comparison with alkali Also size of H + is very small compared to

that of alkali metal ion. H forms stable hydride only with strongly electropositive metals due to smaller value of its electron affinity (72.8 kJ mol -1 ) .

  • In view of the anomalous behavior of hydrogen, it is difficult to assign any definite position to it in the periodic Hence it is customary to place it in group I (Along with alkali metals) as well as in group VII (Along with halogens).
  • Discovery and occurrence : It was discovered by Henry Cavendish in 1766. Its name hydrogen was proposed by Lavoisier. Hydrogen is the 9th most abundant element in the earth’s
  • Preparation of Dihydrogen : Dihydrogen can be prepared by the following methods,
    • Laboratory method : In the laboratory, dihydrogen can be prepared by the action of

H 2 SO4 on granulated Zinc, Zn + H 2 SO4 (dil.) ® ZnSO4 + H 2

  • Industrial method
  • By the electrolysis of water : The hydrogen prepared by this method is highly Dihydrogen is


collected at cathode.

2H 2 O(l) ¾¾Elec¾trol¾y¾sis ® 2H 2(g) + O2(g)


  • Hydrocarbon steam process : H 2 is prepared by the action of steam on e.g.




CH4  + H 2 O ¾¾117¾0¾K ® CO + 3H 2 (Steam)


  • Bosch process :

H 2  + CO+ H 2 O ¾¾7¾73K¾® CO2  + 2H 2


water gas


Fe2O3 , Cr2O3


  • Lane’s process : H 2

is prepared by passing alternate currents of steam and water gas over red


hot iron. The method consists of two stages,


Oxidation stage :


Iron filings

+ 4 H 2 O ¾¾102¾5-1¾075¾K  ®


Fe3 O4

Magnetic oxide of iron

+ 4 H 2 + 161KJ


Reduction stage :

2Fe3 O4  + 4CO + 4 H 2  ® 6Fe + 4CO2  + 4 H 2 O

water gas


  • Physical properties of dihydrogen : It is a colourless, tasteless and odourless It is slightly soluble in water. It is highly combustible. The Physical constants of atomic hydrogen are,

Atomic radius (pm) – 37 ; Ionic radius of H ion (pm) – 210; Ionisation energy (kJ mol -1 ) – 1312;

Electron affinity (kJ mol -1 ) –72.8; Electronegativity – 2.1.

  • Chemical properties of dihydrogen : Dihydrogen is quite stable and dissociates into


hydrogen atoms only when heated above 2000 K,

H 2  ¾¾200¾0¾K ® H + H . Its bond dissociation energy is


very high,

H 2 ® H + H ;

DH = 435.9 kJ mol -1 . Due to its high bond dissociation energy, it is not very


reactive. However, it combines with many elements or compounds.

  • Action with metals : To forms corresponding hydrides. 2Na + H 2 ¾¾He¾at ® 2NaH ;

Ca + H 2  ¾¾He¾at ® CaH 2 .

With transition metals (elements of d – block) such as Pd, Ni, Pt etc. dihydrogen forms interstitial hydrides in which the small molecules of dihydrogen occupy the interstitial sites in the crystal lattices of these hydrides. As a result of formation of interstitial hydrides, these metals adsorb large volume of hydrogen on their surface. This property of adsorption of a gas by a metal is called occlusion. The occluded hydrogen can be liberated from the metals by strong heating.


  • Reaction with Non-metals :

2H 2

  • O2

¾¾970¾K ® 2H

2O ;

N 2 + 3H

Fe, Mo

2  ¾¾750¾K, P¾ress¾u¾re ®

2NH 3


H 2 + F2

¾¾Da¾rk ® 2HF ;

H 2 + Cl


2  ¾¾673¾K, P¾res¾su¾re ®



H 2 + Br2 ® 2HBr ;

H 2  + I 2  ¾¾673¾K ® 2HI



The reactivity of halogen towards dihydrogen decreases as,

F2 > Cl2 > Br2 > I 2


As a result,

F2 reacts in dark,

Cl2 in the presence of sunlight,


reacts only upon heating while the


reaction with

I 2 occurs in the presence of a catalyst.


  • Reaction with unsaturated hydrocarbons :

as ethylene and acetylene to give saturated hydrocarbons.

H 2 reacts with unsaturated hydrocarbons such


H 2 C = CH 2 + H 2  ¾¾Ni o¾r Pt¾or P¾d ® CH3  – CH3  ;  HC º CH+ 2H 2  ¾¾Ni o¾r Pt¾or P¾d ® CH3  – CH3



473 K



473 K



This reaction is used in the hydrogenation or hardening of oils. The vegetable oils such as groundnut oil or cotton-seed oil are unsaturated in nature because they contain at least one double bond



in their molecules. Dihydrogen is passed through the oils at about 473 K in the presence of catalyst to form solid fats. The vegetable ghee such as Dalda, Rath, etc. are usually prepared by this process.


Vegetable oil+ H 2  ¾¾473¾K  ® Fat



(6)  Uses of Dihydrogen



  • As a reducing agent, (ii) In the hydrogenation of vegetable oils, (iii) As a rocket fuel in the form of


liquid H 2

  • In the manufacture of synthetic petrol, (v) In the preparation of many compounds such as


NH 3 , CH3OH, Urea etc, (vi) It is used in the oxy-hydrogen torch for welding if temperature around 2500°C is required. It is also used in atomic hydrogen torch for welding purposes in which temperature of the order of 4000°C is required.

Different forms of hydrogen

  • Atomic hydrogen : It is obtained by the dissociation of hydrogen molecules. The atomic hydrogen is stable only for a fraction of a second and is extremely reactive. It is obtained by passing dihydrogen gas at atmospheric pressure through an electric

arc struck between two tungsten rods.

The electric arc maintains a temperature around 4000 – 4500°C. As the molecules of dihydrogen gas pass through the electric arc, these absorb energy and get dissociated into atoms as

H 2 (g) ¾¾Elec¾t¾ric ® 2H(g) : DH  = 435.90KJ mol -1


This arrangement is also called atomic hydrogen torch.

  • Nascent hydrogen : The hydrogen gas prepared in the reaction mixture in contact with the substance with which it has to react, is called nascent hydrogen. It is also called newly born hydrogen. It is more reactive than ordinary For example, if ordinary hydrogen is passed through acidified



(pink in colour), its colour is not discharged. On the other hand, if zinc pieces are added to the


same solution, bubbles of hydrogen rise through the solution and the colour is discharged due to the


reduction on


by nascent hydrogen.


KMnO4  +

H 2


+ H 2SO4 ® No Re action ;

Zn + H 2SO4 ® ZnSO4 +

2[H] ´ 5

Nascent hydrogen


2KMnO4 + 3H 2 SO4  + 10H ® KSO4  + 2MnSO4  + 8H 2 O

  • Ortho and para hydrogen : A molecule of dihydrogen contains two The nuclei of both the atoms in each molecule of dihydrogen are spinning. Depending upon the direction of the spin of the nuclei, the hydrogen is of two types,
  • Molecules of hydrogen in which the spins of both the nuclei are in the same directions, called ortho
  • Molecules of hydrogen in which the spins of both the nuclei are in the opposite directions, called para



Ordinary dihydrogen is an equilibrium mixture of ortho and para hydrogen. Ortho hydrogen ⇌ Para hydrogen. The amount of ortho and para hydrogen varies with temperature as,

  • At 0°K, hydrogen contains mainly para hydrogen which is more
  • At the temperature of liquefaction of air, the ratio of ortho and para hydrogen is 1:1.
  • At the room temperature, the ratio of ortho to para hydrogen is 3:1.
  • Even at very high temperatures, the ratio of ortho to para hydrogen can never be more than 3:1.

Thus, it has been possible to get pure para hydrogen by cooling ordinary hydrogen gas to a very low temperature (close to 20 K) but it is never possible to get a sample of hydrogen containing more than 75% of ortho hydrogen. i.e., Pure ortho hydrogen can not be obtained.

Isotopes of Hydrogen

Isotopes are the different forms of the same element which have the same atomic number but different mass numbers.

Isotopes of hydrogen



Name   Symbol Atomic number Mass number Relative abundance Nature radioactive or non-radioactive
Protium or Hydrogen   1 H or H 1 1 99.985% Non-radioactive
Deuterium   2 H or D 1 2 0.015% Non-radioactive
Tritium   3 H or T 1 3  1015 %  Radioactive




Physical constants of H2 , D2 and T2

Property H2 D2 T2
Molecular mass 2.016 4.028 6.03
Melting point (K) 13.8 18.7 20.63
Boiling point (K) 20.4 23.9 25.0
Heat of fusion (kJ mol -1 ) 0.117 0.197 0.250
Heat of vaporisation (kJ mol -1 ) 0.994 1.126 1.393
Bond energy (kJ mol -1 ) 435.9 443.4 446.9


Water is the oxide of hydrogen. It is an important component of animal and vegetable matter. Water constitutes about 65% of our body. It is the principal constituent of earth’s surface.

  • Structure : Due to the presence of lone pairs, the

geometry of water is distorted and the HOH bond angle is

104.5°, which is less than the normal tetrahedral angle (109.5°) . The geometry of the molecule is regarded as angular or bent. In

water,  each   O H bond   is   polar   because  of   the   high

electronegativity of oxygen (3.5) in comparison to that of hydrogen (2.1). The resultant dipole moment of water molecule is 1.84D.

In ice, each oxygen atom is tetrahedrally surrounded by four hydrogen atoms; two by covalent bonds and two by hydrogen bonds. The resulting structure of ice is open structure having a number of



vacant spaces. Therefore, the density of ice is less than that of water and ice floats over water. It may be noted that water has maximum density (1g cm-3 ) at 4°C.

  • Heavy water : Chemically heavy water is deuterium oxide (D2O). It was discovered by Urey. It has been finding use in nuclear reactors as a moderator because it slows down the fast moving neutrons and therefore, helps in controlling the nuclear fission
  • Physical properties : Water is colourless, odourless and tasteless liquid at ordinary

Some physical constants of H2O and D2O at 298 K

  • Chemical properties : Water shows a versatile chemical behaviour. It behaves as an acid, a base, an oxidant, a reductant and as ligand to
    • Dissociation of water : Water is quite stable and does not dissociate into its elements even at high Pure water has a small but measurable electrical conductivity and it dissociates as,



HO+  + OH ;


KW   = 1.0 ´ 10 -14 mol 2 L2 at 298K


  • Amphoteric nature : Water can act both as an acid and a base and is said to be However, water is neutral towards litmus and its pH is 7.
  • Oxidising and reducing nature : Water can act both as an oxidising and a reducing agent in


its chemical reactions. e.g.

2Na +

2H 2 O

Oxidi sin g agent

® 2NaOH + H 2 ;

2F2 +

2H 2 O

Re ducing agent

® 4 HF + O2


(5)  Hard and Soft water

Water which produces lather with soap solution readily is called soft water. e.g. distilled water, rain water and demineralised water.

Water which does not produce lather with soap solution readily is called hard water. e.g. sea water, river water, well water and tap water.

  • Cause of hardness of water : The hardness of water is due to the presence of bicarbonates, chlorides and sulphates of calcium and



Hard water does not produce lather because the cations (Ca +2and Mg +2 ) present in hard water react


with soap to form insoluble precipitates,

M +2

  • 2C17 H35 COONa ® (C17 H35 COO)2 M + 2Na + ,Where



M = Ca or Mg

From hard water

Sodium stearate(soap)

Metal stearate(PPt.)


Therefore, no lather is produced until all the calcium and magnesium ions are precipitated. This also results into wastage of lot of soap.

  • Type of hardness of water : The hardness of water is of two types,
  • Temporary hardness : This is due to the presence of bicarbonates of calcium and magnesium. It is also called carbonate
  • Permanent hardness : This is due to the presence of chlorides and sulphates of calcium and It is also called non-carbonate hardness.
    • Softening of water : The process of the removal of hardness from water is called softening of
  • Removal of temporary hardness : It can be removed by the following methods,

By boiling : During boiling, the bicarbonates of Ca and Mg decompose into insoluble carbonates and give CO2 . The insoluble carbonates can be removed by filtration.

Ca(HCO3 )2  ¾¾He¾at ® CaCO3 + CO2  + H 2 O ;  Mg(HCO3 )2  ¾¾He¾at ® MgCO3 + CO2  + H 2 O







Clark’s method : This process is used on a commercial scale. In this process, calculated amount of lime [Ca(OH)2 ]is added to temporary hard water.

Ca(HCO3 )2 + Ca(OH)2 ¾¾® 2CaCO3 ¯ +2H 2 O






Mg(HCO3 )2 + Ca(OH2 ) ¾¾® MgCO3 ¯ +CaCO3  ¯ +2H 2 O






  • Removal of permanent hardness : Permanent hardness can be removed by the following methods,

By washing soda method : In this method, water is treated with a calculated amount of washing soda (Na2CO3 ) which converts the chlorides and sulphates of Ca and Mg into their respective carbonates which get precipitated.


CaCl2 + Na2 CO3 ¾¾® CaCO3 + 2NaCl  ;


MgSO4  + Na2 CO3  ¾¾® MgCO3 + Na2 SO4



Permutit method : This is a modern method employed for the softening of hard water. hydrated sodium aluminium silicate (Na 2 Al2 Si2 O8 .xH 2 O) is called permutit. These complex salts are also known as zeolites.

The permutit as loosely packed in a big tank over a layer of coarse sand. Hard water is introduced into the tank from the top. Water reaches the bottom of the tank and then slowly rises through the permutit layer in the tank. The cations present in hard water are exchanged for sodium ions. Therefore this method is also called ion exchange method.




Na2 Z+

Sodium zeolite

Ca +2 (From hard water)

¾¾® CaZ+ 2Na + ;

Cal zeolite

Na2 Z +

Mg +2 (From hard water)



Magnesium zeolite

+ 2Na + , where


Z = Al 2 Si2 O8 .

xH 2 O

Hydrogen peroxide


Hydrogen peroxide (H 2 O2 )

was discovered by French chemist Thenard.


  • Preparation : It is prepared by
    • Laboratory method : In laboratory,


is prepared by Merck’s process. It is prepared by


adding calculated amounts of sodium peroxide to ice cold dilute (20%) solution of

Na2 O2 + H 2 SO4 ¾¾® Na2 SO4 + H 2 O2

H 2 SO4 .


  • Industrial method : On a commercial scale,


can be prepared by the electrolysis of 50%



solution. In a cell, peroxy disulphuric acid is formed at the anode.

2H 2 SO4  ¾¾Elec¾roly¾sis ® H 2 S2 O8 (aq.)+ H 2 (g)

Peroxy disulphuric acid


This is drawn off from the cell and hydrolysed with water to give

H 2 O2 .


H 2 S2 O8  + 2H 2 O ¾¾® 2H 2 SO4  + H 2 O2

The resulting solution is distilled under reduced pressure




gets distilled while


with high boiling point, remains undistilled.


(2)  Physical properties : Pure


is a thick syrupy liquid with pale blue colour. It is more


viscous and dense than water. It is completely miscible with water, alcohol and ether in all proportions.



(3)  Chemical properties

  • Decomposition : Pure


is an unstable liquid and decomposes into water and

O2 either


upon standing or upon heating,

2HO2 ¾¾® 2HO + O2 ;

DH = -196.0 kJ


  • Oxidising nature : It is a powerful oxidising It acts as an oxidising agent in neutral, acidic


or in alkaline medium. e.g.

2KI + H 2O2 ¾¾® 2KOH + I 2

[In neutral medium]


2FeSO4  + H 2 SO4  + H 2 O2 ¾¾® Fe2 (SO4 )3 + 2H 2 O

[In acidic medium]


MnSO4  + H 2 O2 + 2NaOH ¾¾® MnO2 + Na2 SO4  + 2H 2 O

[In alkaline medium]


  • Reducing nature :


has tendency to take up oxygen from strong oxidising agents and


thus, acts as a reducing agent,


even neutral medium.

HO2 + O ¾¾® HO + O2 . It can act as a reducing agent in acidic, basic or

From oxidising agent


In acidic medium,

H 2 O2 ¾¾® 2H + + O2 + 2e


In alkaline medium,

HO2 + 2OH ¾¾® 2HO + O2 + 2e


  • Bleaching action :


acts as a bleaching agent due to the release of nascent oxygen.


HO2 ¾¾® HO + O




Thus, the bleaching action of


is due to oxidation. It oxidises the colouring matter to a


colourless product, Colouring matter +O ® Colour less matter

H 2O2 is used to bleach delicate materials like ivory, silk, wool, leather etc.


(4)   Structure of   H2O2 :


has non-planar structure in which two H-atoms are arranged in two


directions almost perpendicular to each other and to the axis joining the two oxygen atoms. The O O

linkage is called peroxide linkage.


  • Strength of H2O2 : The strength of


is expressed in terms of weight or volume,


  • As weight percentage : The weight percentage of


gives the weight of


in 100 g of


solution. For example, a 40% solution by wt. means 40 g of H 2O2

are present in 100 g of solution.


  • As volume : The strength of

H 2 O2

is commonly expressed as volume. This refers to the volume


of oxygen which a solution of


will give. For example, a “20 volume” of

H 2 O2

means that 1 litre of


this solution will give 20 litres of oxygen at NTP.

(6)  Uses of H2O2

  • It is used as an antichlor in bleaching because it can reduce

Cl2 + H 2 O2 ¾¾® 2HCl + O2

  • It is used for restoring the colour of lead
  • It is used as an antiseptic for washing wounds, teeth and ears under the name

The group 1 of the periodic table contains six elements, namely lithium (Li), sodium (Na), potassium (K), rubidium (Rb), caesium (Cs) and francium (Fr). All these elements are typical metals. These are usually referred to as alkali metals since their hydroxides forms strong bases or alkalies.

Electronic configuration


Elements Discovery Electronic configuration ( ns1 )
3 Li Arfwedson (1817)  1s2 2s1 or [He]2 2s1
11 Na Davy (1807)  1s 2 2s 2 2p6 3s1  or [Ne]10 3s1
19 K Davy (1807)  1s 2 2s 2 2p6 3s 2 3p6 4s1 or [Ar]18 4s1
37 Rb Bunsen (1861)  1s 2 2s 2 2p6 3s 2 3p6 3d10 4s 2 4 p6 5s1 or [Kr]36 5s1
55 Cs Bunsen (1860)  1s 2 2s 2 2p6 3s 2 3p6 3d10 4s 2 4 p6 4d10 5s 2 5p6 6s1 or [Xe]54 6s1
87 Fr Percy (1939)  1s 2 2s 2 2p6 3s 2 3p6 3d10 4s 2 4 p6 4d10 4 f 14 5s 2 5 p6 5d10 6s 2 6 p6 7s1 or  [Rn]86 7s1



Note : ® Francium is radioactive with longest lived isotope minute.

223 Fr

with half life period of only 21


  • Because of similarity in electronic configuration, they exhibit similar properties. A regular gradation in their properties with increase in no. is observed due to increasing size of atoms/ions and the low binding energy of valency electrons.



  • Of all the alkali metals, only sodium and potassium are found in abundance in Francium occurs only in minute quantities as a radioactive decay product.

Physical properties

  • Physical state
    • All are silvery white, soft and light These can be cut with the help of knife. When freshly cut, they have bright lustre which quickly tarnishes due to surface oxidation.
    • These form diamagnetic colourless ions since these ions do not have unpaired electrons, (i.e. M+

has ns0configuration). That is why alkali metal salts are colourless and diamagnetic.

(2)  Atomic and ionic radii

  • The alkali metals have largest atomic and ionic radii than their successive elements of other groups belonging to same
  • The atomic and ionic radii of alkali metals, however, increases down the group due to progressive addition of new energy

No doubt the nuclear charge also increases on moving down the group but the influence of addition of energy shell predominates

  Li Na K Rb Cs Fr
Atomic radius (pm) 152 186 227 248 265 375
Ionic radius of M+ ions (pm)

(3) Density

60 95 133 148 169
  • All are light metals, Li, Na and K have density less than water. Low values of density are because these metals have high atomic volume due to larger atomic On moving down the group the atomic size as well as atomic mass both increase but increase in atomic mass predominates over increase in atomic size or atomic volume and therefore the ratio mass/volume i.e. density gradually increases down the groups
  • The density increases gradually from Li to Cs, Li is lightest known metal among

Li = 0.534, Na = 0.972, K = 0.86, Rb = 1.53 and Cs = 1.87 g/ml at 200C.

  • K is lighter than Na because of its unusually large atomic
  • In solid state, they have body centred cubic

(4)  Melting point and Boiling point

  • All these elements possess low pt and in comparison to other group members.

Li                            Na                         K                             Rb                            Cs

Fr (K)               453.5            370.8           336.2             312.0             301.5 (K)                1620             1154.4         1038.5           961.0             978.0

  • The lattice energy of these atoms in metallic crystal lattice relatively low due to larger atomic size and thus possess low and On moving down the group, the atomic size increases and binding energy of their atoms in crystal lattice decreases which results lowering of pts.



  • Lattice energy decreases from Li to Cs and thus pt and also decrease from Li to Cs.

(5)  Ionisation energy & electropositive or metallic character

  • Due to unpaired lone electron in ns sub-shell as well as due to their larger size, the outermost electron is far from the uncleus, the removal of electron is easier and these low values of ionisation (I.E.)
Li Na K Rb Cs
520 495 418 403 376
7296 4563 3069 2650 2420


  • Ionisation energy of these metal decreases from Li to Ionisation energy




A jump in 2nd ionisation energy (huge difference) can be explained as,



Li :1s 2 2s1

Re movalof

¾¾® Li + :1s 2 2s electron

Re moval of


1s electron

Li 2+ : 1s1


Removal of 1s electrons from Li+ and that too from completely filled configuration requires much more energy and a jump in 2nd ionisation is noticed

  • Lower are ionisation energy values, greater is the tendency to lose ns1 electron to change in M+

ion      (i.e. M ® M+ + e) and therefore stronger is electropositive character.

  • Electropositive character increases from Li to Cs.

Due to their strong electropositive character, they emit electrons even when exposed to light showing

photoelectric effect. This property is responsible for the use of Cs and K in photoelectric cell.

(6)  Oxidation number and valency

  • These elements easily form univalent + ve ion by losing solitary ns1 electron due to low ionisation energy
  • Alkali metals are univalent in nature and form ionic compounds. Lithium salts are, however,
  • Further, the M+ ion acquires the stable noble gas configuration. It requires very high values of energy to pull out another electron from next to outer shell of M+ ion and that is why their second ionisation energy is very Consequently, under ordinary conditions, it is not possible for these metals to form M2+ ion and thus they show +1 oxidation state.
  • Since the electronic configuration of M+ ions do not have unpaired electron and thus alkali metal salts are diamagnetic and Only those alkali metal salts are coloured which have coloured anions

e.g. K2Cr2O7 is orange because of orange coloured Cr2O 2- ion, KMnO4 is violet because of violet coloured

MnO41- ion.

(7)  Hydration of Ions

  • Hydration represents for the dissolution of a substance in water to get adsorb water molecule by weak valency force. Hydration of ions is the exothermic process when ions on dissolution water get
  • The hydration is an exothermic process i.e energy is released during
  • The energy released when 1 mole of an ion in the gaseous state is dissolved in water to get it
(g)                                   (aq)

hydrated is called hydration energy M            + + Aq ® M +     ; DH = – energy.

  • Smaller the cation, greater is the degree of Hydration energy, Li+ > Na+ > K+ > Rb+

> Cs+



  • Li+ being smallest in size has maximum degree of hydration and that is why lithium salts are mostly hydrated, LiCl. 2H2O Also lithuim ion being heavily hydrated, moves very slowly under the influence of electric field and, therefore, is the poorest conductor current among alkali metals ions It may, therefore, be concluded that it is the degree of hydration as well as the size of ion is responsible for the current carried by an

Relative ionic radii                           Cs+ > Rb+ > K+ > Na+ > Li+ Relative hydrated ionic radii              Li+ > Na+ > K+ > Rb+ > Cs+ Relative conducting power                    Cs+ > Rb+ > K+ > Na + > Li+

(8)    Electronegativities

  • These metals are highly electropositive and thereby possess low values of
Li Na K Rb Cs Fr
0.98 0.93 0.82 0.82 0.79


  • Electronegativity of alkali metals decreases down the group as the trend of numerical values of electronegativity given below on Pauling scale



Note : ® Fr being radioactive elements and thus studies on physical properties of this element are limited.

  • Specific heat : It decreases from Li to




Li Na K Rb Cs
Specific heat (Cal/g)

0.941 0.293 0.17 0.08 0.049
  • Conduction power : All are good conductors of heat & electricity, because of loosely held valence

(11)  Standard oxidation potential and reducting properties

  • Since alkali metals easily lose ns1 electron and thus they have high values of oxidation potential



M + aq ® M +     + e

  • The standard oxidation potentials of a alkali metals (in volts) are listed below,


Li Na K Rb Cs
+3.05 +2.71 +2.93 +2.99 +2.99
  • More is oxidation potential, more is the tendency to get oxidized and thus more powerful is reducing nature in aqueous That is why alkali metals liberate H2 from H2O and HCl.

2H 2 O + 2M ® 2MOH + H 2 ;  2HCl + 2M ® 2MCl + H 2

  • However, an examination of ionisation energy for alkali metals reveals that Li should have the minimum tendency to lose electron and thus its reducing nature should be minimum. The greatest reducing nature of Li in medium is accounted due to the maximum hydration energy of Li+ ion. For Lithium


Li(s) ® Li(g) ;

(g)               ( g )

Li      ® Li+     + e;

DH1 = Heat of sublimation, DHs

DH2 = IE1





( g )


® Li


DH3 = – Heat of hydration, DHh





Li(s) + H 2O ® Li + (aq) + e; DH = DH1 + DH 2 + DH 3 = DHs IE1 – DHh

(s)           2                     ( sq)                                               5           1             h

Similarly, for sodium, Na   + H O ® Na +     + e; DH = DH   + IE   – DH

DHh for Li > DHh for Na. Therefore, large negative DH values are observed in case of Li and this explains for more possibility of Li to get itself oxidized or have reducing nature.

  • Characteristic flame colours : The alkali metals and their salts give characteristic colour to Bunsen The flame energy causes and excitation of the outermost electron which on reverting back to its initial position gives out the absorbed energy as visible light. These colour differ from each other Li – crimson, Na–Golden yellow, K – Pale violet , Rb and Cs –violet. These different colours are due to different ionisation energy of alkali metals. The energy released is minimum in the case of Li+ and increases in the order.

Energy released          : Li+ < Na+ < K+ < Rb+ < Cs+

l released                : Li+ > Na+ > K+ > Rb+ > Cs+

Frequency released     : Li+ < Na+ < K+ < Rb+ < Cs+

Chemical properties

  • Occurrence : Alkali metals are very reactive and thus found in combined state Some important ores of alkali metals are given
  • Lithuim : Triphylite, Petalite, lepidolite, Spodumene [LiAl(SiO3)3], Amblygonite [Li(Al F)PO4]
  • Sodium : Chile salt petre (NaNO3), Sodium chloride (NaCl), Sodium sulphate (Na2SO4), Borax (Na2B4O710H2O), Glauber salt (Na2 SO4.10H2O)
  • Potassium : Sylime (KCl), carnallite (KCl.MgCl2.6H2O)and Felspar (K2Al2O3.6SiO2)
  • Rubidium : Lithuim ores Lepidolite, triphylite contains 7 to 3% Rb2 O
  • Caesium : Lepidolite, Pollucite contains 2 to 7% Cs2O
  • Extraction of alkali metals : Alkali metals cannot be extracted by the usual methods for the extraction of metals due to following
  • Alkali metals are strong reducing agents, hence cannot be extracted by reduction of their oxides or other
  • Being highly electropositive in nature, it is not possible to apply the method of displacing them from their salt solutions by any other
  • The aqueous solutions of their salts cannot be used for extraction by electrolytic method because hydrogen ion is discharged at cathode instead of an alkali metal ions as the discharge potentials of alkali metals are However, by using Hg as cathode, alkali metal can be deposited. The alkali metal readily combines with Hg to form an amalgam from which its recovery difficult. The only successful method, therefore, is the electrolysis of their fused salts, usually chlorides. Generally, another metal chloride is added to lower their fussion temperature.


Fused NaCl :

NaCl ¾¾fusi¾on ® Na +  + Cl

Electrolysis : Anode : 2Cl ® Cl2 + 2e

of fused salt : Cathode : 2Na + + 2e ® 2Na


(3)  Alloys Formation

  • The alkali metals form alloys among themselves as well as with other
  • Alkali metals also get dissolved in mercury to form amalgam with evolution of heat and the amalgamation is highly exothermic .



(4)  Formation of oxides and hydroxides

  • These are most reactive metals and have strong affinity for O2 quickly tranish in air due to the formation of a film of their oxides on the These are, therefore, kept under kerosene or paraffin oil to protect them from air,

M + O® MO ¾¾® MO2

Oxide                 Peroxide

  • When burnt air (O2), lithium forms lithium oxide (Li2O) sodium forms sodium peroxide (Na2O2) and other alkali metals form super oxide (Mo2 e. KO2,RbO2 or CsO2)


2Li + 1 O ®

2   2

Li2 O      ; 2Na + O2 ® Na2 O2  ; K + O2 ®           KO2


Lithuim oxide                                                                            Potassium super oxide

The reactivity of alkali metals towards oxygen to form different oxides is due to strong positive field around each alkali metal cation. Li + being smallest, possesses strong positive field and thus combines with small anion O2- to form stable Li2O compound. The Na+ and K+ being relatively larger thus exert less

strong positive field around them and thus reacts with larger oxygen anion i.e, O2 and O1 to form stable

2               2


The monoxide, peroxides and superoxides have O2 and O2 , O1 ions respectively. The structures of



  • · ·· ··

2        2

  • · ··


each are,                   [ · O· ]2 ; [ · OO· ]2– , [ · O· · · O· ]1

  • · ·                  ·                   ·                           ·

´´                      ··           ··

The O2–1 ion has a three electron covalent bond and has one electron unpaired . It is therefore superoxides are paramagnetic and coloured KO2 is light yellow and paramagnetic substance.

  • The oxides of alkali metals and metal itself give strongly alkaline solution in water with evolution of heat


M + H

O ® MOH + 1 H  ;

2                             2   2

DH = ve




Li2O + H 2O ® 2LiOH;

Na2 O2 + 2H 2 O ® 2NaOH + H 2 O2(l) ; 2KO2 + 2H 2 O ® 2KOH + H 2 O2(l) + O2(g) ;

DH = ve

DH = ve

DH = ve


The peroxides and superoxides act as strong oxidising agents due to formation of H2O2

  • The reactivity of alkali metals towards air and water increases from Li to Cs that is why lithium decomposes H2O very slowly at 25oC whereas Na does so vigorously, K reacts producing a flame and Rb, Cs do so


M + H

O ® MOH + 1 H

2                             2   2



  • The basic character of oxides and hydroxides of alkali metals increases from Li to This is due to the increase in ionic character of alkali metal hydroxides down the group which leads to complete dissociation and leads to increase in concentration of OH ions.



(5)  Hydrides

  • These metal combines H to give white crystalline ionic hydrides of the general of the formula
  • The tendency to form their hydrides, basic character and stability decreases from Li to Cs since the electropositive character decreases from Cs to Li.

2M+ H2 ® 2MH ; Reactivity towards H2 is Cs < Rb < K < Na < Li

  • The metal hydrides react with water to give MOH & H2 ; MH + H2 O ® MOH + H2
  • The ionic nature of hydrides increases from Li to Cs because of the fact that hydrogen is present in the these hydrides as H and the smaller cation will produce more polarisation of anion (according to Fajan rule) and will develop more covalent
  • The electrolysis of fused hydrides give H2 at NaH fused Contains Na + and H i.e.,



At cathode: Na+ +e ® Na   ;     At anode:

H ® 1 H   + e

2   2



  • Alkali metals also form hydrides like NaBH4, LiAlH4 which are good reducing

(6)  Carbonates and Bicarbonates

  • The carbonates (M2CO3) & bicarbonates (MHCO3) are highly stable to heat, where M stands for alkali
  • The stability of these salts increases with the increasing electropositive character from Li to Cs. It is therefore Li2 CO3 decompose on heating, Li2CO3® Li2O+CO2
  • Bicarbonates are       decomposed      at       relatively       low       temperature,
3                            2        3          2                 2

2MHCO   ¾¾300¾0¾C ® M  CO   + H  O + CO

  • Both carbonates and bicarbonates are soluble in water to give alkaline solution due to hydrolysis of carbonate ions or bicarbonate

(7)  Halides

  • Alkali metals combine directly with halogens to form ionic halide M + X .
  • The ease with which the alkali metals form halides increases from Li to Cs due to increasing electropositive character from Li to
  • Lithium halides however have more covalent Smaller is the cation, more is deformation of anion and thus more is covalent nature in compound. Also among lithium halides, lithium iodide has maximum covalent nature because of larger anion which is easily deformed by a cation (The Fajan’s rule)

Thus covalent character in lithium halides is, LiI > LiBr > LiCl > LiF

  • These are readily soluble in water. However, lithium fluoride is sparingly soluble. The low solubility of LiF is due to higher forces of attractions among smaller Li+ and smaller F ions (high lattice energy).
  • Halides having ionic nature have high and good conductor of current. The melting points of halides shows the order, NaF > NaCl > NaBr > Nal
  • Halides of potassium, rubidium and caesium have a property of combining with extra halogen atoms forming




KI + I2 ®KI3 ; In KI3(aq) the ions K+ and I are present

(8)  Solubility in liquid NH3

  • These metals dissolve in liquid NH3 to produce blue coloured solution, which conducts electricity to an appreciable
  • With increasing concentration of ammonia, blue colour starts changing to that of metallic copper after which dissolution of alkali metals in NH3
  • The metal atom is converted into ammoniated metal in e. M+ (NH3) and the electron set free combines with NH3 molecule to produce ammonia solvated electron.


Na + (x + y) ® NH 3 [Na(NH 3 )x ]+ +

Ammoniated cation

[e(NH 3 )y ]

Ammoniated electron


  • It is the ammoniated electron which is responsible for blue colour, paramagnetic nature and reducing power of alkali metals in ammonia solution. However, the increased conductance nature of these metals in ammonia is due to presence of ammoniated cation and ammonia solvated
  • The stability of metal-ammonia solution decreases from Li to
  • The blue solution on standing or on heating slowly liberates hydrogen, 2M + 2NH3 ®2MNH2 + H2 . Sodamide (NaNH2) is a waxy solid, used in preparation of number of sodium
  • Nitrates : Nitrates of alkali metals (MNO3) are soluble in water and decompose on heating. LiNO3 decomposes to give NO2 and O2 and rest all give nitrites and

2MNO3 ® 2MNO2 + O2 (except Li) ; 4 LiNO3 ®2Li2O + 4NO2 + O2

(10)  Sulphates

  • Alkali metals’ sulphate have the formula M2SO4 .
  • Except Li2SO4, rest all are soluble
  • These sulphates on fusing with carbon form sulphides, M2SO4 + 4C ® M2S + 4CO
  • The sulphates of alkali metals (except Li) form double salts with the sulphate of the trivalent metals like Fe, Al, Cr The double sulphates crystallize with large number of water molecules as alum. e.g. K2SO4 . Al2 (SO4)3. 24 H2O.

(11)  Reaction with non-metals

  • These have high affinity for non-metals. Except carbon and nitrogen, they directly react with hydrogen, halogens, sulphur, phosphorus to form corresponding compounds on heating.


2Na + H2

¾¾300¾0¾C ® 2NaH;  2K + H ®2KH



2Na + Cl2 ® 2NaCl                ; 2K + Cl2 ® 2KCl

2Na + S ® Na2S            ; 2K + S ® K2S

3Na + P ® Na3P             ; 3K + P ® K3P

  • Li reacts, however directly with carbon and nitrogen to form carbides and nitrides. 2Li + 2C ® LiC2 ; 6Li + 2N2 ® 2 Li3N
  • The nitrides of these metals on reaction with water give NH3. M3N + 3H2O ® 3MOH + NH3



  • Reaction with acidic hydrogen : Alkali metals react with acids and other compounds containing acidic hydrogen (i.e, H atom attached on F,O, N and triply bonded carbon atom, for example, HF, H2O, ROH, RNH2, CHºCH) to liberate H2 .


M + H

O ® MOH + 1 H      ;

2                             2   2

M + HX ® MX + 1 H

2   2




M + ROH ® ROH + 1 H    ;

2   2


® RNHNa + 1 H

2                              2   2



  • Complex ion formation : A metal shows complex formation only when it possesses the following characteristics, (i) Small size (ii) High nuclear charge (iii) Presence of empty orbitals in order to accept electron pair Only Lithium in alkali metals due to small size forms a few complex ions Rest all alkali metals do not possess the tendency to form complex ion.

Anomalous behaviour of Lithium

Anomalous behaviour of lithium is due to extremely low size of lithium its cation On account of small size and high nuclear charge, lithium exerts the greatest polarizing effect out of all alkali metals on negative ion. Consequently lithium ion possesses remarkable tendency towards solvation and develops covalent character in its compounds. Li differs from other alkali metals in the following respects,

  • It is comparatively harder than other alkali
  • It can be melted in dry air without losing its
  • Unlike other alkali metals, lithium is reactive among It can be noticed by the following properties,

(i) It is not affected by air. (ii) It decomposes water very slowly to liberate H2. (iii) It hardly reacts with bromine while other alkali metals react violently.

  • Lithium is the only alkali metal which directly reacts with N2.
  • Lithium when heated in NH3 forms imide, Li2 NH while other metals form amides, MNH2.
  • When burnt in air,, lithium form Li2O sodium form Na2O and Na2O2 other alkali metals form monoxide, peroxide and
  • Li2O is less basic and less soluble in water than other alkali
  • LiOH is weaker base than NaOH or KOH and decomposes on

2LiOH ¾¾D ® Li2O + H 2O

  • LiHCO3 is liquid while other metal bicarbonates are


  • Only Li2CO3 decomposes on heating not decompose on

Li2 CO3  ¾¾he¾at ® Li2 O + CO2 . Na2CO3, K2CO3 etc. do


  • LiNO3 and other alkali metal nitrates give different products on heating 4LiNO3 = 2Li2O + 4NO2 + O2 ; 2NaNO3 = 2NaNO2 + O2
  • LiCl and LiNO3 are soluble in alcohol and other organic These salts of other alkali metals are, however, insoluble in organic solvents.



  • LiCl is deliquescent while NaCl, KBr etc. are not. Lithium chloride crystals contain two molecules of water of crystallisation (LiCl. 2H2O). Crystals of NaCl KBr, KI etc do not conation water of
  • Li2SO4 does not form alums like other alkali
  • Li reacts with water slowly at room temperature Na reacts vigorously Reaction with Rb and Cs is violent.
  • Li reacts with Br2 Reaction of other alkali metals with Br2 is fast.
  • Li2 CO3 Li2C2O4, LiF , Li3PO4 are the only alkali metal salts which are insoluble or sparingly soluble in

Diagonal Relationship of Li with Mg

Due to its small size lithium differs from other alkali metals but resembles with Mg as its size is closer to Mg Its resemblance with Mg is known as diagonal relationship. Generally the periodic properties show either increasing or decreasing trend along the group and vice versa along the period which brought the diagonally situated elements to closer values. Following are the characteristic to be noted.

Period           Group I         Group II

  • Li Be
  • Na Mg
  • Both Li and Mg are harder and higher than the other metals of their groups.
  • Due to covalent nature, chlorides of both Li and Mg are deliquescent and soluble in alcohol and pyridine while chlorides of other alkali metals are not
  • Fluorides, phosphates of li and Mg are sparingly soluble in water whereas those of other alkali metals are soluble in
  • Carbonates of Li and Mg decompose on heating and liberate CO2 Carbonates of other alkali metals are stable towards heat and decomposed only on

Li2CO3 ® Li2O + CO2 ;  Mg CO3 ® MgO + CO2

  • Hydroxides and nitrates of both Li and Mg decompose on heating to give Hydroxides of both Li and Mg are weak alkali.

4 LiNO3 ® 2Li2O + 4NO2 + O2 ; 2Mg(NO3)2 ®2MgO + 4NO2 + O2

2LiOH ® Li2O + H2O ; Mg(OH)2 ® MgO + H2O

Hydroxides of other alkali metals are stable towards heat while their nitrates give O2 and nitrite.

2KNO3 ® 2KNO2 + O2

  • Both Li and Mg combine directly with N2 to give nitrides Li3 N and Mg3 N2. Other alkali metals combine at high temperature, 6Li + N2 ® 2Li3N; 3Mg + N2 ® Mg3 N2. Both the nitrides are decomposed by water to give NH3

Li3N + 3H2O ® 3LiOH + NH3 ; Mg3N2 + 6H2o ® 3Mg(OH)2+ 2NH3

  • Bicarbonates of Li and Mg are more soluble in water than carbonates whereas carbonates whereas carbonates of alkali metals are more
  • Both Li and Mg combine with carbon on



2Li + 2C ® Li2C2 ; Mg + 2C ® Mg C2

  • The periodic properties of Li and Mg are quite comparable


  Li Mg
Electronegativity 1.0 1.2
Atomic radii 1.23 1.36
Ionic radii 0.60(Li+) 0.65(Mg+2)
Atomic volume 12.97 c.c 13.97 c.c
  • Both have high polarizing Polarizing Power = Ionic charge / (ionic radius)2.
  • Lithium and Mg Form only monooxide on heating in 4Li + O2 ® 2 Li2O ; 2Mg + O2 ® 2 MgO
  • Li2SO4 Like MgSO4 does not form
  • The bicarbonates of Li and Mg do not exist in solid state, they exist in solution
  • Alkyls of Li and Mg (R.Li and MgX) are soluble in organic solvent.
  • Lithium chloride and MgCl2 both are deliquescent and separate out from their aqueous solutions as hydrated crystals, 2H2O and MgCl2 . 2H2O.

Sodium and its compounds


  • Ores of sodium : NaCl (common salt), salt), borax (sodium tetraborate or sodium borate,

NaNO3 (chile salt petre),

Na2 B4 O7 .10H 2 O) .

Na2 SO4 .10H 2 O (Glauber’s


  • Extraction of sodium : It is manufactured by the electrolysis of fused sodium chloride in the presence of CaCl2 and KF using graphite anode and iron This process is called Down process.

NaCl Na + + Cl .


At cathode :

Na + + e ® Na ; At anode :

Cl ® Cl + e ;

Cl + Cl ® Cl2 ­




Note : ® Sodium cannot be extracted from aqueous NaCl because



H2O / H2

(–0.83V) is more


than E 0 Na + / Na (–2.71V).




Cl2 .

  • Anode and cathode are separated by means of a wire gauze to prevent the reaction between Na


(3)  Compound of sodium

Sodium hydroxide (Caustic soda), NaOH


  • Preparation
  • Gossage process :

Na2 CO3

(10% solution)

  • Ca(OH)2 ® 2NaOH ¯ +CaCO3


  • Electrolytic method : Caustic soda is manufactured by the electrolysis of a concentrated solution of NaCl.


At anode: Cl discharged; At cathode:

Na +



  • Castner – Kellener cell (Mercury cathode process) : NaOH obtained by electrolysis of

solution of brine. The cell comprises of rectangular iron tank divided into three compartments.






Outer compartment – Brine solution is electrolysed ; Central compartment – 2% NaOH solution

H 2

  • Properties : White crystalline solid, highly soluble in water, It is only sparingly soluble in


  • Reaction with salt :

FeCl3 + 3NaOH ®

Fe(OH)3    ¯ +

(Insoluble hydroxide)



HgCl 2 + 2NaOH ® 2NaCl

  • Hg(OH)2 ®


H 2O + HgO ¯





AgNO3  + 2NaOH ® 2NaNO3  + 2AgOH ®

Ag 2 O ¯ + H 2 O Brown




Note : ®

Zn, Al, Sb, Pb, Sn

and As forms insoluble hydroxide which dissolve in excess of NaOH


(amphoteric hydroxide).

  • NH 4 Cl + NaOH ¾¾he¾at ® NaCl + NH 3 ­ +H 2 O


  • Reaction with halogens :

X 2 + 2NaOH

(cold) ®

NaX + NaXO + H 2O

sod. hypohalite


3X 2 + 6 NaOH (hot) ® 5 NaX + NaXO3 + 3H 2 O ;

(Sod. halate)

(X = Cl, Br, I)


  • Reaction with metals: Weakly electropositive metals like

Zn + 2NaOH ® Na2 ZnO2 + H 2 ­

Zn, Al

and Sn etc.


  • Reaction with sand, SiO2 :

2NaOH + SiO2 ®

Na2 SiO3

Sod. silicate (glass)

  • H 2 O


  • Reaction with CO:

NaOH + CO ¾¾150¾20¾0o¾C ® HCOONa


510 atm             Sod. formate

Note : ® NaOH breaks down the proteins of the skin flesh to a pasty mass, therefore it is commonly known as caustic soda.

  • Uses : In the manufacturing of sodium metal, soap, rayon, paper, dyes and For mercuring cotton to make cloth unshrinkable and reagent in lab.

Sodium carbonate or washing soda, Na 2CO3


  • Preparation : Solvay process : In this process, brine (NaCl) ,

NH 3

and CO2 are the raw


NH 3 + CO2 + H 2 O ® NH 4 HCO3


NH 4


  • NaCl


30 C

¾¾ ¾®


¯ + NH 4 Cl




250 C

3 ¾¾¾¾®



  • H 2

O + CO2


2NH 4 Cl + Ca(OH)2 ® CaCl2  + 2H 2 O + 2NH 3

slaked lime

Note : ® CaCl 2 so formed in the above reaction is a by product of solvay process.


  • Properties : (a)

Na2 CO3 .10H 2 O ¾¾dry¾a¾ir ® Na2 CO3 .H 2 O+ 9H 2 O


(decahydrate)                      (Monohydrate)




Na2 CO3 . H 2 O ¾¾®

Na2 CO3  ®

It does not decompose on further heating even to redness ( 853o C)


  • It is soluble in water with considerable evolution of

Na2 CO3 + H 2 O ® H 2 CO3  + 2Na + + 2OH

Weak acid

  • It is readily decomposed by acids with the evolution of CO2
  • Na2 CO3 + H 2 O + CO2 ® 2NaHCO3
  • Uses : In textile and petroleum refining, Manufacturing of glass, NaOH soap powders

Sodium peroxide (Na2O2)

  • Preparation : It is manufactured by heating sodium metal on aluminium trays in air (free from CO2 )

2Na + O2 (air) ¾¾® Na2 O2

  • Properties : (a) When pure it is colourless. The faint yellow colour of commercial product is due to presence of small amount of superoxide (NaO2 ).


  • On coming with moist air it become white due to formation of NaOH and

Na2 CO3 .


2Na2 O2 + 2H 2 O ® 4 NaOH + O2  ;  2NaOH + CO2  ® Na2 CO3  + H 2 O

  • It is powerful oxidising agent. It oxidises Cr (III) hydroxide to sodium chromate, Mn (II) to sodium manganate and sulphides to
  • Uses : As a bleaching agent and it is used for the purification of air in confined spaces such as


submarines because it can combine with CO2 to give

2CO2 + 2Na2 O2 ® 2Na2 CO3  + O2 .

Na2 CO3

and oxygen,



The group 2 of the periodic table consists of six metallic elements. These are beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr), barium (Ba) and radium (Ra). These (except Be) are known as alkaline earth metals as their oxides are alkaline and occur in earth crust.

Electronic configuration


Element Electronic configurations ( ns 2 )
4 Be  1s 2 2s 2 or [He] 2s 2
12 Mg  1s 2 2s 2 2p6 3s 2 or [Ne]3s 2
20 Ca  1s 2 2s 2 2p6 3s 2 3p6 4s 2 or [Ar]4s 2
38 Sr  1s 2 2s 2 2p6 3s 2 3p6 3d10 4s 2 4 p6 5s 2 or [Kr]5s 2
56 Ba  1s 2 2s 2 2p6 3s 2 3p6 3d10 4s 2 4 p6 4d10 5s 2 5 p6 6s 2 or [Xe]6s 2
88 Ra  1s 2 2s 2 2p6 3s 2 3p6 3d10 4s 2 4 p6 4d10 4 f 14 5s 2 5p6 5d10 6s 2 6p6 7s 2 or [Rn]7s 2

Note : ® Radium was discovered in the ore pitch blende by madam Curie. It is radioactive in nature.

Physical properties



  • Physical state : All are greyish-white, light, malleable and ductile metals with metallic Their hardness progressively decrease with increase in atomic number. Althought these are fairly soft but relatively harder than alkali metals.

(2)  Atomic and ionic radii

  • The atomic and ionic radii of alkaline earth metals also increase down the group due to progressive addition of new energy shells like alkali
  Be Mg Ca Sr Ba Ra
Atomic radius (pm) 112 160 197 215 222
Ionic radius of M2+ ion (pm) 31 65 99 113 135 140
  • The atomic radii of alkaline earth metals are however smaller than their corresponding alkali metal of the same This is due to the fact that alkaline earth metals possess a higher nuclear charge than alkali metals which more effectively pulls the orbit electrons towards the nucleus causing a decrease in size.

(3)  Density

  • Density decreases slightly upto Ca after which it increases. The decrease in density from Be to Ca might be due to less packing of atoms in solid lattice of Mg and
Be Mg Ca Sr Ba      Ra
1.84 1.74 1.55 2.54 3.75    6.00
  • The alkaline earth metals are more denser, heavier and harder than alkali metal. The higher density of alkaline earth metals is due to their smaller atomic size and strong intermetallic bonds which provide a more close packing in crystal lattice as compared to alkali

(4)  Melting point and Boiling point

  • Melting points and boiling points of alkaline earth metals do not show any regular
  Be Mg Ca Sr Ba Ra (K) 1560 920 1112 1041 1000 973 (K) 2770 1378 1767 1654 1413
  • The values are, however, more than alkali This might due to close packing of atoms in crystal lattice in alkaline earth metals.

(5)  Ionisation energy and electropositive or metallic character

  • Since the atomic size decreases along the period and the nuclear charge increases and thus the electrons are more tightly held towards nucleus. It is therefore alkaline earth metals have higher ionisation energy in comparison to alkali metals but lower ionisation energies in comparison to p-block
  • The ionisation energy of alkaline earth metals decreases from Be to
  Be Mg Ca Sr Ba Ra
First ionisation energy (k J mol-1) 899 737 590 549 503 509
Second ionisation energy (kJ mol-1) 1757 1450 1146 1064 965 979
  • The higher values of second ionisation energy is due to the fact that removal of one electron from the valence shell, the remaining electrons are more tightly held in which nucleus of cation and thus more energy is required to pull one more electron from monovalent



  • No doubt first ionisation energy of alkaline earth metals are higher than alkali metals but a closer look on 2nd ionisation energy of alkali metals and alkaline earth metals reveals that 2nd ionisation energy of alkali metals are more


Li       Be

1st ionisation energy (kJ mol-1)               520     899 2nd ionisation energy (kJ mol-1)  7296   1757

This may be explained as, Li : 1s2, 2s1 ¾¾rem¾ova¾l of ¾2s ® Li +: 1s2




¾¾rem¾ova¾l of¾1s ®




Li2+ : 1s1


Be : 1s2 , 2s2

¾¾rem¾ova¾l of ¾2s ® Be+ : 1s2, 2s1


¾¾rem¾ova¾l of ¾2s ® Be2+ : 1s2



The removal of 2nd electron from alkali metals takes place from 1s sub shell which are more closer to nucleus and exert more nuclear charge to hold up 1 s electron core, whereas removal of 2nd electron from alkaline earth metals takes from 2s sub shell. More closer are shells to the nucleus, more tightly are held electrons with nucleus and thus more energy is required to remove the electron.

  • All these possess strong electropositive character which increases from Be to
  • These have less electropositive character than alkali metals as the later have low values of ionisation

(6)  Oxidation number and valency

  • The IE1 of the these metals are much lower than IE1 and thus it appears that these metals should form univalent ion rather than divalent ions but in actual practice, all these give bivalent This is due to the fact that M2+ ion possesses a higher degree of hydration or M2+ ions are extensively hydrated to form [M(H2O)x]2+, a hydrated ion. This involves a large amount of energy evolution which counter balances the higher value of second ionisation energy.

M ®M2+ ,      DH = IE1 + E2

M2+ + xH2O ® [M(H2O)x]2+; DH = – hydration energy.

  • The tendency of these metals to exist as divalent cation can thus be accounted as,
  • Divalent cation of these metals possess noble gas or stable
  • The formation of divalent cation lattice leads to evolution of energy due to strong lattice structure of divalent cation which easily compensates for the higher values of second ionisation energy of these
  • The higher heats of hydration of divalent cation which accounts for the existence of the divalent ions of these metals in solution

(7)  Hydration of ions

  • The hydration energies of alkaline earth metals divalent cation are much more than the hydration energy of monovalent

Mg+              Mg2+

Hydration energy or Heat of hydration (kJ mol-1)         353           1906

The abnormally higher values of heat of hydration for divalent cations of alkaline earth metals are responsible for their divalent nature. MgCl2 formation occurs with more amount of heat evolution and thus MgCl2 is more stable.

  • The hydration energies of M2+ ion decreases with increase in ionic

Be2+      Mg2+     Ca2+      Sr2+      Ba2+



Heat of hydration kJ mol-1 2382 1906   1651   1484   1275

  • Heat of hydration are larger than alkali metals ions and thus alkaline earth metals compounds are more extensively hydrated than those of alkali metals g MgCl2 and CaCl2 exists as Mg Cl2 .6H2O and CaCl2. 6H2O which NaCl and KCl do not form such hydrates.
  • The ionic mobility, therefore, increases from Ba2+ to Ba2+, as the size of hydrated ion decreases.

(8)  Electronegativities

  • The electronegativities of alkaline earth metals are also small but are higher than alkali
  • Electronegativity decreases from Be to Ba as shown below,


  Be Mg Ca Sr Ba
Electronegativity 1.57 1.31 1.00 0.95 0.89
  • Conduction power : Good conductor of heat and

(10)  Standard oxidation potential and reducing properties

  • The standard oxidation potential (in volts) are,


Be Mg Ca Sr Ba
1.69 2.35 2.87 2.89 2.90
  • All these metals possess tendency to lose two electrons to give M2+ ion and are used as reducing
  • The reducing character increases from Be to Ba, however, these are less powerful reducing agent than alkali
  • Beryllium having relatively lower oxidation potential and thus does not liberate H2 from

(11)   Characteristic flame colours

  • The characteristic flame colour shown are : Ca-brick red; Sr –crimson ; Ba-apple green and Ra-
  • Alkaline earth metals except Be and Mg produce characteristic colour to flame due to easy excitation of electrons to higher energy
  • Be and Mg atoms due to their small size, bind their electrons more strongly (because of higher effective nuclear charge) Hence these requires high excitation energy and are not excited by the energy of flame with the result that no flame colour is shown by

Chemical properties

  • Occurrence : These are found mainly in combined state such as oxides, carbonates and sulphates Mg and Ca are found in abundance in nature. Be is not very abundant, Sr and Ba are less Ra is rare element. Some important ores of alkaline earth metals are given below,
  • Baryllium : Beryl (3BeO.Al2O3.6SiO2); Phenacite (Be2SiO4)
  • Magnesium : Magnesite (MgCO3); Dolomite (CaCO3. MgCO3); Epsomite(MgSO 7H2O); Carnallite (MgCl2.KCl. 6H2O); Asbestos [CaMg3(SiO3)4]
  • Calcium : Limestone (CaCO3); Gypsum : (CaSO4.2H2O), Anhydrite (CaSO4); Fluorapatite [(3Ca3(PO4)CaF2)] Phosphorite rock [Ca3(PO4)2]
  • Barium : Barytes (BaSO4) ; witherite (BaCO3)



  • Radium : Pitch blende (U3O8); (Ra in traces); other radium rich minerals are carnotite [K2UO2)] (VO4)2 8H2O and antamite [Ca(UO2)2]

(2)  Extraction of alkaline earth metals

  • Be and Mg are obtained by reducing their oxides carbon,

BeO + C ® Be + CO ; MgO + C ® Mg + CO

  • The extraction of alkaline earth metals can also be made by the reduction of their oxides by alkali metals or by electrolysing their fused
  • Alloy formation : These dissolve in mercury and form

(4)  Formation of oxides and hydroxides

  • The elements (except Ba and Ra) when burnt in air give oxides of ionic nature M2+O2- which are crystalline in nature . Ba and Ra however give peroxide. The tendency to form higher oxides increases from Be to

2M + O2 ® 2MO       (M is Be, Mg or Ca ) 2M + O2 ® MO2               (M is Ba or Sr)

  • Their less reactivity than the alkali metals is evident by the fact that they are slowly oxidized on exposure to air, However the reactivity of these metals towards oxygen increases on moving down the
  • The oxides of these metals are very stable due to high lattice
  • The oxides of the metal (except BeO and MgO) dissolve in water to form basic hydroxides and evolve a large amount of BeO and MgO possess high lattice energy and thus insoluble in water.
  • BeO dissolves both in acid and alkalies to give salts i.e. BeO possesses amphoteric nature. BeO + 2NaOH ® Na2BeO2 + H2O ; BeO + 2HCl ® BeCl2 + H2O

Sod. beryllate                                                            Beryllium chloride


  • The basic nature of oxides of alkaline earth metals increases from Be to Ra as the electropositive Character increases from Be to
  • The tendency of these metal to react with water increases with increase in electropositive character e. Be to Ra.
  • Reaction of Be with water is not certain, magnesium reacts only with hot water, while other metals react with cold water but slowly and less energetically than alkali
  • The inertness of Be and Mg towards water is due to the formation of protective , thin layer of hydroxide on the surface of the
  • The basic nature of hydroxides increase from Be to It is because of increase in ionic radius down the group which results in a decrease in strength of M –O bond in M –(OH)2 to show more dissociation of hydroxides and greater basic character.
  • The solubility of hydroxides of alkaline earth metals is relatively less than their corresponding alkali metal hydroxides Furthermore, the solubility of hydroxides of alkaline earth metals increases from Be to Be (OH)2 and Mg (OH)2 are almost insoluble, Ca (OH)2 (often called lime water) is sparingly soluble whereas Sr(OH) 2 and Ba (OH)2 (often called baryta water ) are more soluble.



The trend of the solubility of these hydroxides depends on the values of lattice energy and hydration energy of these hydroxides. The magnitude of hydration energy remains almost same whereas lattice energy decreases appreciably down the group leading to more –Ve values for DH solution down the group.

DH solution = DH lattice energy + DH hydration energy

More negative is DH solution more is solubility of compounds.

  • The basic character of oxides and hydroxides of alkaline earth metals is lesser than their corresponding alkali metal oxides and
  • Aqueous solution of lime water [Ca(OH)2] or baryta water [Ba(OH)]2 are used to qualitative identification and quantative estimation of carbon dioxide, as both of them gives white precipitate with CO2 due to formation of insoluble CaCO3 or BaCO3

Ca(OH)2 + CO2 ® CaCO3 + H2O ; Ba(OH)2 + CO2 ® BaCO3 + H2O

(white ppt)                                                                   (white ppt)


Note : ® SO2 also give white ppt of CaSO3 and BaSO3 on passing through lime water or baryta water However on passing CO2 in excess, the white turbidity of insoluble carbonates dissolve to give a clear solution again due to the formation of soluble bicarbonates,

CaCO3 ® H2O + CO2 ® Ca(HCO3)2

(5)  Hydrides

  • Except Be, all alkaline earth metals form hydrides (MH2) on heating directly with H2 . M+ H2 ®





  • BeH2 is prepared by the action of Li Al H4 On BeCl2 ; 2BeCl2 + LiAlH4 ® 2BeH2 + LiCl +


  • BeH2 and MgH2 are covalent while other hydrides are
  • The ionic hydrides of Ca, Sr, Ba liberate H2 at anode and metal at




Ca2+ + 2H


Anode : 2H ®H2 + 2e Cathode : Ca2+ + 2e ®Ca

  • The stability of hydrides decreases from Be to Ba.
  • The hydrides having higher reactivity for water, dissolves readily and produce hydrogen gas.

CaH2(s) + 2H2O ®Ca(OH) 2 + 2H2­

(6)  Carbonates and Bicarbonates

  • All these metal carbonates (MCO3) are insoluble in neutral medium but soluble in acid medium . These are precipitated by the addition of alkali metal or ammonium carbonate solution to the solution of these

(NH4)2 CO3 + CaCl2 ®2NH4Cl + CaCO3  ; Na2CO3 + BaCl2 ®2NaCl + BaCO3

  • Alkaline earth metal carbonates are obtained as white precipitates when calculated amount of carbon dioxide is passed through the solution of the alkaline metal

M(OH)2 (aq) + CO2 (g) ®MCO3(s) + H2O(l)



and sodium or ammonium carbonate is added to the solution of the alkaline earth metal salt such as CaCl2.

CaCl2 (aq) + Na2CO3 (aq) ® CaCO3(s) +2 NaCl(aq)

  • Solubility of carbonates of these metals also decreases downward in the group due to the decrease of hydration energy as the lattice energy remains almost unchanged as in case of

(vi) The carbonates of these metals decompose on heating to give the oxides, the temperature of decomposition increasing from Be to Ba. Beryllium carbonate is unstable.

MCO3  ¾¾He¾at ® MO + CO2





(7)  Halides

  • The alkaline earth metals combine directly with halogens at appropriate temperatures forming halides, MX2. These halides can also be prepared by the action of halogen acids (HX) on metals, metal oxides, hydroxides and

M + 2HX ® MX2 + H2  ;  MO + 2HX ®MX2 + H2O

M(OH)2 + 2HX® MX2 +2H2O  ;  MCO3 + 2HX ® MX2 + CO2 + H2O

Beryllium chloride is however, conveniently obtained from oxide

BeO + C + Cl2  ¾¾870¾10¾70¾K ® BeCl2  + CO.

  • BeCl2 is essentially covalent, the chlorides MgCl2, CaCl2 , SrCl2 and BaCl2 are ionic; the ionic character increases as the size of the metal ion The evidence is provided by the following facts,
  • Beryllium chloride is relatively low melting and volatile whereas BaCl2 has high melting and
  • Beryllium chloride is soluble in organic
    • The halides of the members of this group are soluble in water and produce neutral solutions from which the hydrates such : MgCl2 6H2O, CaCl2.6H2BaCl2 2H2O can be crystallised. The tendency to form hydrated halides decreases with increasing size of the metal ions.
    • BeCl2 is readily hydrolysed with water to form acid solution, BeCl2 + 2H2O ®Be (OH)2 +
    • The fluorides are relatively less soluble than the chlorides due to high lattice energies. Except BeCl2 and MgCl2 the chlorides of alkaline earth metals impart characteristic colours to

CaCl2                    SrCl2                    BaCl2

Brick red colour      Crimson colour      Grassy green colour

Structure of BeCl2 In the solid phase polymeric chain structure with three centre 2 electron bonding with Be-Cl-Be bridged structure is shown below,



202 PM

Cl                                                                                                                  Cl


Be     98   82

Cl                                                                                                                   Cl



In the vapour phase it tends to form a chloro-bridged dimer which dissociates into the linear triatomic monomer at high temperature at nearly 1200 K.

  • Solubility in liquid ammonia : Like alkali metals, alkaline earth metals also dissolve in liquid ammonia to form coloured solutions When such a solution is evaporated, hexammoniate, M(NH3)6 is

(9)  Nitrides

  • All the alkaline earth metals directs combine with N2 give nitrides, M3N2.
  • The ease of formation of nitrides however decreases from Be to
  • These nitrides are hydrolysed water to liberate NH3, M3N2 + 6H2O ®3M(OH)2 + 2NH3

(10)  Sulphates

  • All these form sulphate of the type M SO4 by the action of H2 SO4 on metals, their oxides, carbonates or

M + H2SO4 ® MSO4 + H2  ; MO + H2SO4 ® MSO4 + H2O

MCO3+ H2SO4 ® MSO4 + H2O+CO2 ; M(OH)2 + H2SO4 ® MSO4 + 2H2O

  • The solubility of sulphates in water decreases on moving down the group BeSO4 and MgSO4 are fairly soluble in water while BaSO4 is completely insoluble. This is due to increases in lattice energy of sulphates down the group which predominates over hydration

(ii) Sulphate are quite stable to heat however reduced to sulphide on heating with carbon.

MSO4 + 2C ®MS+2CO2

  • Action with carbon : Alkaline metals (except Be, Mg) when heated with carbon form carbides of the type MC2 These carbides are also called acetylides as on hydrolysis they evolve

MC2 + 2H2O®M(OH) 2 + C2H2

  • Action with sulphur and phosphorus : Alkaline earth metals directly combine with sulphur and phosphorus when heated to form sulphides of the type MS and phosphides of the type M3P2

M + S ® MS ; 3M + 2P ® M3P2

Sulphides on hydrolysis liberate H2S while phosphides on hydrolysis evolve phosphine.

MS + dil. acid ® H2S ; M3P2 + dil. acid ® PH3 Sulphides are phosphorescent and are decomposed by water

2MS + 2H2O® M(OH) 2 + M(HS)2

  • Nitrates : Nitrates of these metals are soluble in water On heating they decompose into their corresponding oxides with evolution of a mixture of nitrogen dioxide and


M(NO3 )2

® MO + 2NO2

æ 1 ö

+              O

ç 2 ÷ 2


è   ø

(14)  Formation of complexes

  • Tendency to show complex ion formation depends upon smaller size, high nuclear charge and vacant orbitals to accept Since alkaline metals too do not possess these characteristics and thus are unable to form complex ion.
  • However, Be2+ on account of smaller size forms many complex such as (BeF3)1-, (BeF4)2-.

Anomalous behaviour of Beryllium



Beryllium differs from rest of the alkaline earth metals on account of its small atomic size, high electronegativity Be2+ exerts high polarizing effect on anions and thus produces covalent nature in its compounds. Following are some noteworthy difference of Be from other alkaline earth metals,

  • Be is lightest alkaline earth
  • Be possesses higher pt. and than other group members.
  • BeO is amphoteric in nature whereas oxides of other group members are strong
  • It is not easily effected by dry air and does not decompose water at ordinary
  • BeSO4 is soluble in
  • Be and Mg carbonates are not precipitated by (NH4)2 CO3 in presence of NH4
  • Be and Mg salts do not impart colour to
  • Be does not form peroxide like other alkaline earth
  • It does not evolve hydrogen so readily from acids as other alkaline earth metals do
  • It has strong tendency to form complex
  • Be3N2 is volatile whereas nitrides of other alkaline earth metals are non-volatile.
  • It’s salts can never have more than four molecules of water of crystallization as it has only four available orbitals in its valence
  • Berylium carbide reacts water to give methane whereas magnesium carbide and calcium carbide give propyne and acetylene

Be2C+4H2O®2Be(OH)2 + CH4 ; Mg2C3 + 4H2O ® 2Mg(OH)2 + C3H6


CaC2 + 2H2O ® Ca(OH)2 + C2H4

Diagonal relationship of Be with Al

Due to its small size Be differs from other earth alkaline earth metals but resembles in many of its properties with Al on account of diagonal relationship.

  • Be2+ and Al3+ have almost same and smaller size and thus favour for covalent
  • Both these form covalent compounds having low pt and soluble in organic solvent.
  • Both have same value of electronegativity (i.e.1.5).
  • The standard P of these elements are quite close to each other ; Be2+=1.69 volts and Al3+=

1.70 volts.

  • Both become passive on treating with HNO3 in cold.
  • Both form many stable complexes g. (BeF3), (AlH4).
  • Like BeO, Al2O3 is amphoteric in Also both are high m. pt. solids.

Al2O3 + 2NaOH ® 2NaAlO2 + H2O ; Al2O3 + 6HCl ® 2AlCl3 + 3H2O

  • Be and Al both react with NaOH to liberate H2 forming beryllates and

Be + 2NaOH®Na2BeO2+H2 ; 2Al + 6NaOH® 2Na3AlO3 + 3H2

  • Be2 C and Al4C3 both give CH4 on treating with

Be2C+ 2H2O®CH4 + 2BeO ; Al4C3 + 6H2O®3CH4 + 2Al2O3

  • Both occur together in nature in beryl ore, Al2O3. 6SiO2.
  • Unlike other alkaline earths but like aluminium, beryllium is not easily attacked by air (Also Mg is not attacked by air)
  • Both Be and Al react very slowly with HCl to liberate H2.
  • Both Be and Al form polymeric covalent hydrides while hydrides of other alkaline earth are


  • Both BeCl2 and AlCl3 are prepared is similar



BeO+ C+ Cl2 ®BeCl2 + CO ; Al2O3 + 3C +3Cl2 ® 2AlCl3 + 3CO

  • Both BeCl2 and AlCl3 are soluble in organic solvents and act as catalyst in Friedel –Crafts
  • Both Be (OH)2 and Al (OH) 3 are amphoteric whereas hydroxides of other alkaline earths are strong


  • The salts of Be and Al are extensively
  • BeCl2 and AlCl3 both have a bridged polymeric
  • Be and Al both form fluoro complex ions [BeF4]2- and [AlF6]3- in solution state whereas other

members of 2nd group do not form such complexes.

Difference between alkali metals and alkaline earth metals


Properties Alkaline earth metals Alkali metals
Electronic configuration Two electrons are present in the valency One  electron   is   present   in   the
  shell. The configuration is ns2 valency shell.  The  configuration  is
Valency Bivalent Monovalent
Electropositive nature Less electropositive More electropositive
Carbonates Insoluble in water. Decompose Soluble in water. Do not decompose
  On heating on heating (Li2CO3 is an exception).
Hydroxides Weak bases, less soluble and decompose Strong bases,   highly   soluble   and
  on heating stable towards heat.
Bicarbonates These are not known in free state. These are known is solid state.
Action of carbon Exist only in solution Do not directly combine with carbon.
  Directly combine with carbon and form  
Action of nitrogen Directly combine with nitrogen and form Do      not     directly     combine     with
  nitrides nitrogen.
Nitrates Decompose on heating evolving a mixture Decompose on heating evolving Only
  of NO2 and oxygen oxygen
Hydration of compounds The compounds are extensively hydrated. The compounds are less hydrated
  MgCl2. 6H2O,CaCl2, 6H2O, BaCl2, 2H2O

are hydrated chlorides

NaCl, KCl, RbCl form non-hydrated chlorides.
Solubility of salts Sulphates,           phosphates,          fluorides, Sulphates,      phosphates,     fluorides,
  chromates, oxalates etc. are insoluble in chromates, oxalates, etc. are soluble
  water in water
Physical properties Comparatively     harder.      High      melting Soft.           Low          melting         points,
  points. Diamagnetic Paramagnetic

Magnesium and its compounds


  • Ores of magnesium : Magnesite

(MgCO3 ),


(MgCO3 .CaCO3 ), Epsomite (epsom



(MgSO4 .7H 2O)


(MgCl2 . KCl. 6H 2O)


(CaMg 3 (SiO3 )4 ),



(Mg 2 (Si2 O5 )2 . Mg(OH)2 ) .



  • Extraction of magnesium : It is prepared by the electrolysis of fused magnesium chloride which is obtained from carnallite and

(3)  Compounds of magnesium


  • Magnesia (MgO) : It is used as magnesia It is a mixture of MgO and called Sorel’s cement.

MgCl2 . It is also


  • Magnesium hydroxide : It aqueous suspension is used in Medicine as an antacid. Its medicinal name is milk of
  • Magnesium sulphate or Epsom salt (MgSO4 .7H 2 O) : It is isomorphous with ZnSO4 .7H 2 O.

It is used as a purgative in medicine, as a mordant in dyeing and as a stimulant to increase the secretion of bile.

  • Magnesium chloride (MgCl2 .6H 2 O): It is a deliquescent Hydrated salt on heating in air


undergoes partial hydrolysis.

MgCl2 . 6H 2 O ¾¾He¾at ® Mg(OH)Cl + HCl + 5H 2 O .


Calcium and its compounds


  • Ores of calcium : Lime stone or marble or chalk

(CaCO3 ),


(CaSO4 . 2H 2O),



(CaCO3 . MgCO3 ),


(CaF2 ),


Ca3 (PO4 )2 .    Calcium phosphate is a


constituent of bones and teeth.

  • Manufacture : It is manufactured by the electrolysis of a molten mixture of calcium chloride containing some calcium fluoride. Calcium chloride in turn is obtained as a by product of the solvay

(3)  Compounds of calcium

  • Calcium oxide or Quick lime or Burnt lime (CaO) : It’s aqueous suspension is known as slaked

CaO + H 2 O ¾¾hiss¾ing¾so¾un¾d ® Ca(OH)2 + Heat,

slaked lime

When exposed to oxy-hydrogen flame, it starts emitting light called lime light.

Note : ® CaO is used as basic flux, for removing hardness of water, as a drying agent (for

NH3 gas) for preparing mortar (CaO+ sand +water).


  • Calcium chloride (CaCl 2 .6H 2 O) : Fused

CaCl 2

is a good dessicant (drying agent). It can’t be


used to dry alcohol or ammonia as it forms additional products with them.

  • Calcium carbonate (CaCO3) : Ca(OH)2 + CO2 ® CaCO3 + H 2 O .


Note : ® It is insoluble in water but dissolves in the presence of


due to the formation of


calcium bicarbonate.

CaCO3  + H 2 O + CO2  ® Ca(HCO3 )2


  • It is a constituent of protective shells of marine