Shubhang Garodia

IMPORTANT COMPOUNDS AND THEIR FORMULA

1. Baryta water : Ba(OH)2 solution

2. Baking soda : NaHCO3

3. Bleaching powder : CaOCl2

4. Brine : NaCl solution

5. Caustic soda : NaOH

6. Caustic potash : KOH

7. Dry ice : Solid CO2

8. Fusion mixutre : Na2CO3 + K2CO3

9. Fluid magnesia : 12% solution of Mg(HCO3)2

10. Glauber salt : Na2SO4 . 10H2O

11. Hydrolith : CaH2

12. Lime (quick lime or burnt lime) : CaO

13. Mortar : Slaked lime + sand (1 : 3 in water)

14. Micro cosmic salt : NaNH4.HPO4.4H2O

15. Milk of magnesia : Paste of Mg(OH)2 in water (Antacid)

16. Magnesia : MgO

17. Nitrolim : CaCN2 + C (a fertilizer)

18. Nessler's reagent : (K2HgI4 + KOH) aqueous solution

19. Indian saltpetre : KNO3

20. Pearl ash (Potash) : K2CO3

21. Plaster of paris : CaSO4 . ½ H2O or 2CaSO4 . H2O

22. Sorel's cement (Magnesia cement) : MgCl2 . 5MgO . XH2O

23. Soda - lime : NaOH + CaO

24. Soda ash : Na2CO3 (anhydrous)

25. Slaked lime : Ca(OH)2

26. Tincal (Borax) : Na2B4O7 . 10H2O

27. Washing soda : Na2CO3 . 10H2O

28. Salt cake : Na2SO4 (anhydrous)

Shubhang Garodia physical proper ties of s-Block Elements

ALKALI METAL ALKALINE EARTH METAL

Physical state

l One electron in outermost shell & General formula ns1 l Two electrons in outer most shell & General formula ns2.
l Francium is radioactive element. l Radium is radioactive element.
l All are silvery white l All are greyish white.
l Light soft, malleable and ductile metals with metallic lustre. l These metals are harder than alkali metals.
l Alkali metals are paramagnetic, diamagnetic and l These are diamagnetic and colourless in form of ions
colourless in form of ions. or in metal states.

Atomic size

l Biggest in their respective period l Smaller than IA group elements, since extra charge
on nucleus attracts the electron cloud.
(except noble gas element)
l Size increases gradually from Be to Ra
l Size increases from Li to Fr due to addition of an
Be < Mg < Ca < Sr < Ba
extra shell.
Li < Na < K < Rb < Cs < Fr In s-block elements
IA IIA Be is the smallest, Cs is the biggest
Li Be
Na Mg
K Ca
 
Rb Sr
 
Cs Ba

Softness

l Alkali metals are soft because of - l These metals are slightly harder than IA group
(a) Large atomic size because of -
(b) BCC crystal structure (HCP in Li) (a) Smaller atomic size
(c) Loose packing (68% packing efficiency) (b) FCC, HCP crystal structures
(d) Weak metallic bond (c) Packing capacity 74%
(d) Stronger metallic bond due to presence of
l Cs is the softest metal in s-block
two electrons in valence shell.
l Be is the hardest metal in s-block.

1 1
Atomic size   softness 
strength of metallic bond Melting & Boiling point

Melting point and Boiling point

l Weak interatomic bonds are due to their large atomic l Metallic bond is stronger than IA group due to smaller
radii and presence of only one valence electron hence atomic size and two electrons in valence shell hence
melting point and boiling point are low. melting point and boiling point are higher.
l Decreasing order of melting point and boiling point is l Decreasing order of melting point and boiling point is
Li > Na > K > Rb > Cs Be > Ca > Sr > Ba > Mg
l
Shubhang Garodia
With the increase in the size of metal atom, the
repulsion of the non-bonding electrons increases and
l Melting point and Boiling point of Ca, Sr and Ba is
higher than Mg because of presence of d-orbitals in
the outer most shell, which forms stronger metallic
therefore melting point and boiling point decreases
bond.
from Li to Cs.
Melting & Boiling point  Strength of metallic bond

Ionisation potential (I.P.)

l First ionisation potential is higher than IA group
l First ionisation potential (I.P.) is very less because of
because of smaller atomic size and completely filled
bigger atomic size and only one electron in outer
s-orbital (stability)
most shell.
l Second ionisation potential is very high because of l Second ionisation potential is lesser than IA group.
achieving inert gas configuration. l Decreasing order of ionisation potential –
l Decreasing order of ionisation potential -
Be > Mg > Ca > Sr > Ba
Li > Na > K > Rb > Cs
(Ist and 2nd ionisation potential difference < 11eV)
(Ist and 2nd ionisation potential difference > 16eV)
Oxidation state

l The alkali metals shows only + 1 oxidation state. l Alkaline earth metal shows +2. Oxidation state

Electro positive character or metallic character

l Their atomic size is smaller than IA group so these are

l Electropositivity  1/Ionisation Potential
lesser electro positive than IA group. Electropositivity
Due to their larger size electron can easily be removed increases from Be to Ba
to form M+ ion. Electro positive property increases
from Li to Cs.

Density
l Density increases from Be to Ba
l (D = M /V)
As (D = M /V)
In a group atomic volume also increase along with
Exception : (i) Density of Be is higher than Ca and
atomic weight but atomic weight increases more than
Mg because of less volume in comparision to its mass
atomic volume, so density increases from Li to Cs
(ii) Density of Mg is higher than Ca, (reason – As in IA
Exception : Density of Na is higher than K as empty
group.) Increasing order of density -
d-orbitals are present in K. Maximum capacity of M
shell is of 18 electrons but it contains only 8 electron, Ca < Mg < Be < Sr < Ba
which decreases its density. Increasing order of density K < Na < Ca < Mg
Li < K < Na < Rb < Cs
Conductivity
l These are also good conductor of heat and electricity
l Due to the presence of loosely held valence electrons
due to presence of two free electrons.
which are free to move in a metal structure, these
elements are good conductor of heat and electricity. Conductvity of IA < IIA
l
Shubhang Garodia
Alkali metals and their salts gives characteristic
Flame test
l Be and Mg atoms, due to small size, bind their
colour to bunsen flame. The flame energy causes electrons more strongly, so are not excited to higher
an excitation of the outer most electron which on level, hence no flame test.
dropping back to ground state emits absorbed
l Other elements gives characteristic colour to flame
energy as visible light Ex.
Ca-Brick red Sr-dark red
Li-Crimson red Na-Golden yellow K-Violet
Rb-Red violet Cs-Blue Ba-green
Solubility in liquid ammonia
l All the alkali metals dissolves in NH3 (liq.) and produces l Only Ca, Sr and Ba gives blue solution of ammoniated
electron.
blue solution.
l Be and Mg are small in size and have high ionisation
l This blue solution conducts electricity and possesses potential so do not dissolves in liquid NH3.
strong reducing power, due to the presence of l Dark blue colour of solution becomes fade if it allowed
ammoniated electrons. to stand for a long time, it is because of metal amide
Na(s) + (x+y) NH3 [Na(NH3)x]+ + [e(NH3)y]– formation.
ammoniated electron l Blue colour of solution disappears on addition of
ammonium salt, due to NH3 formation.
l This dilute solution is paramagnetic in nature.
NH4+ + NH2  2NH3
Photo electric effect
l Atomic size of K, Rb and Cs is quite large, so their l These elements do not show this property as their
ionisation potential is very low atomic size is small hence ionisation potential is higher
l Due to very low ionisation potential their valence shell than IA group.
electrons gets excited even by absorbing visible light.
That's why Cs is used in photo cells.
Standard oxidation potential
l All the alkali metals have high +ve values of standard l They have lower values of standard oxidation
oxidation potential (tendency of releasing electrons potential due to their small size.
in water or self ionic solutions) l Increasing order of standard oxidation potential is -
l So these are good reducing agent, having upper most
Be < Mg < Ca < Sr < Ba
positions in the electro chemical series.
l Tendendy of loosing electron increases
l Li has highest standerd oxidation potential (+3.05 eV)
due to its high hydration energy. Such that it converts
into Li+ ion by loosing one electron.
Order of standard oxidation potential of s - block element
Li > K > Ba > Sr > Ca > Na > Mg > Be
Hydration energy  Charge density on ion

Hydration energy (Heat of hydration)

l Alkali metals salts are generally soluble in water due
to hydration of cations by water molecules. l Due to smaller ionic size and higher charge density
l Smaller the cation, greater is the degree of its their hydration energy is high.
hydration.
l Li+ Na+ K+ Rb+ Cs+ l Its decreasing order is
Be+2 > Mg+2 > Ca+2 > Sr+2 > Ba+2
– Degree of hydration decreasing
– Hydration energy decreasing l Hydration energy  1/cation size
– Hydrated ion size decreasing
– Ionic conductance increasing
Shubhang Garodia Complex formation tendency

l Only those elements can form complex compounds

l Less tendency to form complex compound, but due
which have
to small size of cations Be and Mg forms complex
(a) Small cation size compounds like

(BeF4)–2, Be4O(CH3COO)6
(b) High charge density
Mg – Chlorophyll
(c) Vacant d-orbital to accept electrons.

l Only Li+ can form complex compound, due to its

small size rest alkali metals have very less tendency
to form complex compounds.

Chemical properties of s-block elements

Reactivity

l These elements are very reactive, so do not found in l Less reactive than alkali metals.
free state in nature. Be < Mg < Ca < Sr < Ba
Reactivity  1/Ionisation potential l BeNo reaction even with hot water
order of reactivity – Li < Na < K < Rb < Cs
Mg  reacts with hot water
l Li is stable reacts slowly with steam. While Rb and
Ca, Sr, Ba  Reacts with cold water.
Cs reacts even with cold water.

Reaction with air

l Alkali metals gets turnish in air due to the formation l Except Be, these metals are easily tarnished in air, as
of oxide at their surface hence they are kept in a Layer of oxide is formed on the surface.
kerosene or paraffin oil. l Barium in powdered form, burst into flame on
exposure to air.
l These elements reacts with moist air to form
carbonates l In moist air, except Be all the elements converts into
carbonates.
4Na + O2  2Na2O
l In dry air Be and Mg gives nitride and oxide both
Na2O + H2O 2NaOH while other gives only oxides.

(moist)

2NaOH + CO2 Na2CO3 + H2O

(in air)

In dry air only Li gives nitride and oxide both while

other elements gives only oxides.
 Shubhang Garodia Reaction with oxygen
l Alkaline earth metals reacts with O2 to form 'MO'
Oxide ion [O2–] :
l Li forms only Li2O (Lithium oxide).
type oxides
(M = Be, Mg, Ca, Sr, Ba)
Peroxide [O2]—2 :
l Na reacts with O2 to form peroxide (Na2O2). l But Ca, Sr and Ba due to low ionisation potential
and more reactivity, forms MO2 (peroxides) at low
temperature.
Ex. CaO2, SrO2, BaO2
l Peroxides are coloured due to Lattice deffect.
Super oxide [O2–] :
l K, Rb and Cs forms MO2 type oxides (super oxides) l BeO shows amphoteric property.
in excess of O2. So super oxides are paramagnetic MgO  weak base
and coloured.
O2 O2 O2
CaO, SrO & BaO  Strong base
M MO MO MO2
2
Oxide
2 2
perioxide super
l Basic properties increases from Be to Ba
oxide
(Li2O) (Na2O2) (KO2, RbO2, CsO2)
l Their stability order is –
Normaloxide > Peroxide > Superoxide

Reaction with hydrogen

l Alkali metals combine with H2 forming ionic hydrides l Except Be all the alkaline metals forms MH2 type
hydrides, (MgH2, CaH2, SrH2, BaH2) on heating
2M + H2  2MH
directly with H2
l LiH is covalent hydride while others are ionic.
l BeH2 is prepared by action of BeCl2 with LiAlH4
l Hydrides of alkali metals are attacked by water to
give back 2BeCl2 + LiAlH4  2BeH2 + LiCl + AlCl3
LiH, NaH, KH, RbH, CsH



Thermal stability decrease l BeH2 and MgH2 are covalent, other are ionic.
 Basic property increases

l H2, are used as reducing agents l Be and Mg have tendency of polymerisation.

MH + H2O  MOH + H2
l electrolysis of fused MH gives H2 at anode.
Reaction with water

l Alkali metals react vigorously with water forming l These metals reacts slowly with water gives H2 and
hydroxides with the liberation of H2. metals hydroxides.
2M + 2H2O  2MOH + H2
M + 2H2O  M(OH)2 + H2
l Reactivity with water increases from Li to Cs.
Li  least reactive towards water l Be does not reacts with water
Na  reacts vigorously l Mg reacts only with hot water
K  reacts producing a flame
l Ca, Sr, Ba reacts with cold water but not as
Rb, Cs  reacts explosively.
energetically as alkali metals.
l These metals also reacts with alcohol gives alkoxide
and H2. order of reactivity Ba > Sr > Ca > Mg > Be
2Li + 2C2H5OH  2C2H5OLi + H2 l from Be(OH)2 to Ba(OH)2 basic property and stability
l Monoxides gives strongly alkaline solution with water
increases.
M2O + H2O  2MOH
 Shubhang Garodia Halides

l Alkali metals reacts directly with halogen to form MX l Alkaline metals reacts with X (Halogen) to form MX2.
Ex. (BeCl2, MgCl2, CaCl2 etc.)
(M – alkali metal, X – Halide ion)
l Ionic nature of MX2 increases from BeCl2 to BaCl2
l Ionic properties of MX increases from LiCl to CsCl
l Ba burns in contact with Cl2
l LiCl is covalent in nature (due to polarisation of Cl–
ion by small Li+ ion). hence it hydrolyses with water l Hydrolytic nature of these halides decreases from
while rest are ionic so do not hydrolyse. BeCl2 to BaCl2

l K, Rb and Cs halides reacts with more halogens to l BeCl2 and MgCl2 are covalent in nature. Order of
gives polyhalides. ionic nature –

on
BeCl2 < MgCl2 < CaCl2 < SrCl2 < BaCl2
KI  I2  KI3 
ionisation
 K   I3
Solubility in water
CsBr + Br2  CsBr3  Cs +
+ Br 3
–
BeCl2 > MgCl2 > CaCl2 > SrCl2 > BaCl2

Carbonates

l All the alkali metals forms M2CO3 type carbonates. l All the alkaline metals forms MCO3 type carbonates.
l Except Li2CO3, all the carbonates are stable towards
l Except BeCO3, all the carbonates are stable towards
heat
heat

Li2CO3 Li2O + CO2 
BeCO3   BeO + CO2
l Thermal stability of carbonates  1/ (Ionic potential)
l Order of decreasing stability -
Order of stability is – BaCO3 > SrCO3 > CaCO3 > MgCO3 > BeCO3
Cs2CO3 > Rb2CO3 > K2CO3 > Na2CO3 > Li2CO3

Nitrates

l Alkali metals forms MNO3 type nitrates (M – alkali l Alkaline earth metals forms M(NO3)2 type nitrates.
metal) (M –Alkaline earth metal).
l Stability increases from LiNO3 to CsNO3. LiNO3 l Stability increases from Be(NO3)2 to Ba(NO3)2 but these
decompoes into Lithium oxide & NO2 on heating. are less stable than IA group, due to smaller atomic size.

4LiNO3   2Li2O + 4NO2 + O2 l All alkaline metals nitrates on heating gives oxides
Oxide
and NO2 + O2
l Other nitrates, on heating to give nitrite and oxygen.

M(NO3)2  Oxides + NO2 + O2

MNO 3  2MNO 2  O2

Nitrite l Be(NO3)2 forms a layer of BeO on its surface so
reaction stops.

Nitrides

l Only Li reacts directly with N2 to form nitride which l Only Be and Mg burns in N 2 to give M3N2 (Be3N2,
gives NH3 on reacting with water. Mg3N2)
6Li + N2  2Li3N Be3N2 + 6H2O 3Be(OH)2 + 2NH3
Li3N + 3H2O  3LiOH + NH3
Mg3N2 + 6H2O  3Mg(OH)2 + 2NH3
Shubhang Garodia Formation of amalgam

l Alkali metals gives amalgam with Hg.

l These metals reacts with other metals to give mixed
metals (alloys) l Shows same properties.

Sulphates
l Alkali metals forms M2SO4 type sulphates. l Alkaline earth metals forms MSO4 type sulphates.
l All alkali metal sulphates are ionic. Ionic properties l Ionic nature of alkaline metal sulphate is increases
increases from Li to Cs.
from Be to Ba
Li2SO4 < Na2SO4 < K2SO4 < Rb2SO4 < Cs2SO4
BeSO4 < MgSO4 < CaSO4 < SrSO4 < BaSO4
l Li2SO4 Least soluble in water.
l Solubility decreases from BeSO4 to BaSO4 as Be+2 and
l These sulphates on burning with C forms sulphides Mg+2 are of small size so their hydration energy is high
Hydration Energy > Lattice energy.
M2SO4 + 4C  M2S + 4CO
l Order of solubility –
l Except lithium, sulphates of IA group reacts with
BeSO4 > MgSO4 > CaSO4 > SrSO4 > BaSO4
sulphates of trivalent metals like Fe+3, Cr+3, Al+3 etc.
gives double salts called alum. l Order of thermal stability –
I III BeSO4, MgSO4, CaSO4, SrSO4, BaSO4

M2SO4.M2(SO4)3.24H2O – Ionic nature increases

–Thermal stability increases

Reaction with acids

l Reacts vigorously with acids. l Freely reacts with acids and displaces hydrogen
2M + H2SO4 M2SO4 + H2 M + 2HCl MCl2 + H2

Reaction with dry NH3

2Li + NH3  Li2NH (Lithimide) l On increasing metal ion concentration solution
converts into bronze colour due to cluster formation
2Na + 2NH32NaNH2 + H2
of metal ions.
(Sodamide)
 Shubhang Garodia CHEMISTRY OF LIGHTER ELEMENTS

Sodium (Na), Natrium

(I) Extraction : Down's Process
By Electrolysis of fused NaCl + CaCl2 + NaF
At cathode (Iron Vessel) : Na+ +e– Na(s)
–
At Anode (Graphite) : 2Cl Cl2 + 2e–
l (CaCl2 + NaF) is used to lower Melting point (8000C) of NaCl to about 6000C.
l Aqueous sodium chloride cannot be used for preparing sodium by electrolysis. Because instead of metallic
sodium, hydrogen gas will be liberated at cathode.
(II) Properties :
(a) It is a crystalline soft metal.
(b) Highly reactive, so kept in kerosene.
(c) Na dissolves in liquid NH3 to giveblue solution.
(III) Uses
(i) In the preparation of sodium amalgam (used as reducing agent)
(ii) In sodium vapour lamp, which emits monochromatic yellow light.
(iii) As heat transfer medium in nuclear reactors.

(1) Sodium chloride NaCl

(I) Occurrence : Sea water is the main source and also found in salt lakes.
(II) Preparation :
Evaporation
(a) Sea water NaCl(2.7 – 2.9%) by solar heat crude NaCl

(b) It contains impurities – Na2SO4, MgCl2, CaCl2 etc.

(c) Insoluble impurities removed by filtration.
HCl gas passed
(d) Filtrate Pure NaCl precipitation (Common ion effect)
_ _
HCl Cl NaCl Na+ + Cl
–
Ionic product of [Na+] [Cl ] > solubility product of NaCl hence it precipitates out.
(e) MgCl2 and CaCl2 are more soluble in water so left in solution.
(III) Properties :
(a) Table salt is slightly hygroscopic due to the presence of magnesium and calcium chlorides in small amounts.
(b) Reaction with AgNO3
NaCl + AgNO3  NaNO3 + AgCl(white ppt.)

(c) 4NaCl + K2Cr2O7 + 5H2SO4 


 4NaHSO4 + K2SO4 + 2CrO2Cl2 + 3H2O
(orange red)
(VI) Uses
(a) As a preservative for pickles, meat and fish.
(b) For making freezing mixture with Ice.
 Shubhang Garodia
(2)
(I)
Sodium Hydroxide (NaOH), Caustic Soda
Preparation : By electrolysis of NaCl.
Castner – Kellner Cell : (Hg – Cathode Process)
_
Electrolite (Brine) NaCl Na+ + Cl
On electrolysis –
At Cathode (Hg)
Na+ + e–  Na. and Na + Hg  Na.Hg (amalgum)
At anode (Graphite)
–
2Cl  Cl2(g) + 2e– and 2Na.Hg + 2H2O  2NaOH + H2 + 2Hg
(II) Properties :
(a) It is deliquescent white crystalline solid.
(b) It absorbs CO2 from air forming Na2CO3.
(c) NaOH is strong base
SiO2
Na2SiO3 + H2O
NaOH Al2O3
2NaAlO2 + H2O

(d) Reaction with non metals : no reaction with H2, N2 and C

B
Na3BO3 (sodium borate)
Si
NaOH Na2SiO3 (sodium silicate)
P4
NaHPO
2 2 + PH3

(sodium hypo phosphite)

(e) Reaction with halogens :
cold/dil
NaX + NaOX (Sodium hypo halite)
NaOH + X2 Hot/conc.
NaX + NaOX3 (Sodium Halate)
(Halogens)

(f) Reaction with Metal :

Alkali metal
No reaction
2Al + 2H2O
2NaAlO2 (Sodium meta aluminate) + H2
B
Na3BeO2 (Sodium Berrylate) + H2
NaOH Zn
Na2ZnO2 (Sodium zincate) + H2
Sn
Na2SnO2 (Sodiumstannate) + H2
Pb
Na2PbO2 (Sodiumplumbite) + H2
(g) The hydroxides of aluminium, zinc, lead and tin, however, dissolve in excess of sodium hydroxide giving
clear solution which can also be obtained when these metals are acted upon by the concentrated solution
of sodium hydroxide.

Zn(OH)2 + 2OH– [Zn(OH)4]2–

–
Al(OH)3 + 3OH 3–
[Al(OH)6]
Zincate ion Aluminate ion
 Shubhang Garodia
(h) Reaction with ZnCl2 or ZnSO4
(a) ZnCl2 + 2NaOH Zn(OH)2 + 2NaCl


(b) Zn(OH)2 + 2NaOH Na2 [Zn(OH)4]
(Soluble complex )

(III) Uses
(a) In the manufacture of soap, rayon, dyes, paper and drugs.
(b) In petroleum refining.
(3) Sodium Bicarbonate or Baking soda (NaHCO3)
(I) Preparation : Solvay process (Commercial Scale)
CaCO3 CaO + CO2 (In brine saturated with NH3, CO2 is passed)
NH3 + H2O + CO2 NH4HCO3
NaCl + NH4HCO3 NH4Cl + NaHCO3
2NH4Cl + CaO CaCl2 + 2NH3 + H2O (Bye-products)
(II) Properties :
(a) Hydrolysis
NaHCO3 + H2O NaOH + H2CO3
(b) Effect of heat (temp. > 100°C)
(Process occurs during preparation of cake)
2NaHCO3 Na2CO3 + H2O + CO2
(c) Reaction with acids – gives CO2
NaHCO3 + HCl NaCl + H2O + CO2
(d) Reaction with base
NaHCO3 + NaOH Na2CO3 + H2O
(III) Uses
(a) In the preparation of baking powder.
(b) In the preparation of effervescent drinks.
(c) In the fire extinguishers.
(d) As antacid medicine (removing acidity)
(4) Sodium carbonate or washing soda (Na2CO3.10H2O)
(I) Occurrence : Na2CO3–Soda ash.
(II) Preparation : Solvay process
(a) Concentrated aqueous solution of NaCl is saturated with NH3.
(b) Current of CO2 passed through the solution.
(c) NaHCO3 precipitated –
NH3 + CO2 + H2O NH4HCO3 NH4+ + HCO3–
–
NaCl Na+ + Cl
[Na+] × [HCO3–] > Ksp of NaHCO3 (so ppt. forms)
2NaHCO3 Na2CO3 + H2O + CO2
(d) Potassium bicarbonate (KHCO3) cannot be prepared by solvay process as it is soluble in water.
(III) Properties :
(a) Efflorescence :
Na2CO3.10H2O when exposed to air it gives out nine out of ten H2O molecules.
Na2CO3.10H2O Na2CO3.H2O + 9H2O
(Monohydrate)
This process is called efflorescence. Hence washing soda losses weight on exposure to air.
 Shubhang Garodia
(b) Hydrolysis :Aqueous solution of Na2CO3 is alkaline in nature due to anionic hydrolysis.
Na2CO3 2Na+ + CO3–2 and CO3–2 + H2O H2CO3 + 2OH
–

(Carbonic acid)
(IV) Uses :
(a) For making fusion mixture (Na2CO3 + K2CO3)
(b) In the manufacture of glass, caustic soda, soap powders etc.
(c) In laundries and softening of water.

(5) Sodium Peroxide Na2O2 :

(I) Preparation :
Sodium peroxide is manufacured by heating sodium metal on aluminium trays in air (free from CO2)
2Na + O2 (air) Na2 O2
(II) Properties :
(a) When pure it is colourless, and the faint yellow colour of the usual product arises from the presence of
a small amount of NaO2.
(b) When it is exposed, it comes in contact with moist air and turns white due to formation of NaOH and
Na2CO3. Thus
Na2O2 + 2H2O 2NaOH + H2O2 and 2NaOH + CO2 Na2CO3 + H2O
(c) Sodium peroxide is a powerful oxidizing agent and oxidizes chromium (III) hydroxide to sodium chromate,
manganese (II) to sodium manganate and sulphides to sulphates.
2Cr(OH)3 + 3O2–2 CrO4–2 + 2OH– + 2H2O
(III) Uses :
(a) Sodium peroxide is widely used as an oxidizing agent yielding in inorganic chemisty; its reaction with
organic compounds are dangerously violent.
(b) Sodium readily combines with carbon dioxide, sodium carbonate and oxygen, it may be used for the
purification of air in confined spaces such as submarines.
(c) It is also used as a bleaching agent because of its oxidizing property.
(d) Sodium peroxide is used in the manufacture of dyes, and many other chemicals such as benzoyl peroxide,
sodium perborate etc.
Magnesium (Mg) :-
(I) Preparation :
(a) From Magnesite or Dolomite : The ore is first calcined to form the oxide
MgCO3 MgO + CO2 and CaCO3.MgCO3 CaO. MgO + 2CO2
(i) From MgO : The oxide is mixed with carbon and heated in a current of chlorine gas
MgO + C + Cl2 MgCl2 + CO
The chloride thus obtained is subjected to electrolysis.
(ii) From CaO.MgO : The mixed oxides (CaO.MgO) obtained from calcination of dolomite
(CaCO3.MgCO3) are reduced by ferrosilicon under reduced pressure above 1273 K.
(b) It is prepared by the electrolysis of fused magnesium chloride.

(II) Properties :
(a) Magnesium burns in air with dazzling light.
2Mg + O2 2MgO
(b) Burning Mg continues to burn in CO2 forming MgO because reducing nature Mg > C
2Mg + CO2 2MgO + C
 Shubhang Garodia
(III) Uses
(a) In preparation of alloy
Electron : 95% Mg + 5% Zn, air craft
Magnalium : 1 – 15% Mg + 85 – 99% Al, used in aeroplanes, balance beams, light instruments.
(b) In photographic flash light.
(c) In preparation of Grignard's reagent.

(1) Magnesium Chloride MgCl2

(I) Occurrence : It is mainly found in sea water and carnallite KCl.MgCl2.6H2O.
(II) Preparation :
(a) By reaction of dil HCl on MgCO3
MgCO3 + 2HCl MgCl2 + CO2 + H2O
(b) MgCl2 is obtained by burning Mg metal in chlorine
Heated
Mg + Cl2 MgCl2
(III) Properties :
On heating MgCl2.6H2O, it gets hydrolysed by its own water of crystallization to an oxy chlorides.

MgCl2.6H2O MgO + 2HCl + 5H2O
(IV) Uses
(a) For preparation of metallic magnesium.
(b) In manufacture of magnesia cement.
(c) Used for dressing cotton threads.

(2) Magnesium Sulphate MgSO4

(I) Occurrence : It occurs naturally as kiserite (MgSO4.H2O) and epsomite (MgSO4.7H2O).

(II) Preparation :
By dissolving magnesite in dil. H2SO4
MgCO3 + H2SO4 MgSO4 + H2O + CO2
(III) Properties : On heating above 2000C
Above
2MgSO4 200°C 2MgO + 2SO2 + O2

Note : It is used in medicine as purgative.

Calcium (Ca) :-
(1) Calcium Oxide (CaO) Quick lime
(I) Preparation : By heating limestones at 8000C.
800°C
CaCO3 CaO + CO2
(II) Properties :
(a) Action of water : CaO + H2O Ca(OH)2
(quick lime) (Slaked lime)
Ca(OH)2 paste in water called milk of lime.
 Shubhang Garodia
(b) Basic Nature :

CaO + SiO2

CaSiO3 and CaO + P4O10

2Ca3(PO4)2
(Calcium silicate)
2000°C
(c) Reaction with carbon : CaO + 3C CaC2 + CO 
(III) Uses
(a) In the manufacture of bleaching powder, cement, glass, calcium carbide etc.
(b) In the purification of sugar
(c) As a drying agent for NH3 and C2H5OH
(d) As basic lining in furnaces
(e) For making Soda lime

(2) Calcium hydroxide Ca(OH)2 Slaked lime

(I) Preparation : By the action of water on quick lime

CaO + H2O Ca(OH)2 + heat
(II) Properties :
(a) Action of CO2 : Lime water turns milky on passing CO2 gas.
Ca(OH)2 + CO2 CaCO3 + H2O
Milkiness
Excess of
CaCO3 CO2
Ca(HCO3)2 (soluble)
(b) Action of Chlorine :
below 35°C
Ca(OH)2 + Cl2 CaOCl2 + H2O
dry Bleaching powder
red heat
2Ca(OH)2 + 2Cl2 2CaCl2 + 2H2O + O2

(III) Uses :
(a) For softening of hard water.
(b) For purification of sugar and Coal gas.
(c) In the manufacture of bleaching powder, Caustic soda and soda lime
(d) In preparation of mortar, plaster and white washing.

(3) Calcium Carbonate (CaCO3) :

Uses :
(a) In the preparation tooth pastes, cosmetics face powder and medicine for indigestion.
(b) In the preparation of Quick lime.
(c) As a building material.
(d) In manufacture of cement, glass, washing soda etc.
(4) Caclium Sulphate CaSO4.2H2O (Gypsum) :
(I) Preparation : CaSO4.2H2O is naturally occuring calcium sulphate. It can be obtained by the action of dil.H2SO4
on a soluble calcium salt below 600C.

CaCl2 + H2SO4 2HCl + CaSO4 

dilute white ppt.
 Shubhang Garodia
(II) Properties :
(a) Action of heat :
120°C 200°C
2CaSO4.2H2O 
–3H 2 O
  (CaSO4)2.H2O   2CaSO4 + H2O

(Plaster of paris) (Anhydride)

(b) It forms an important fertilizer (NH4)2SO4

CaSO4 + 2NH3 + CO2 + H2O  CaCO3  + (NH4)2 SO4
(III) Uses
(a) In the preparation of plaster of paris
(b) Anhydrous CaSO4 used as drying agent.
(c) Anhydrite (CaSO4) is used for manufacture of sulphuric acid, ammonium sulphate.
(5) Plaster of Paris 2CaSO4.H2O
(I) Preparation : It obtained when gypsum is heated at 1200C

2(CaSO4.2H2O) 2CaSO4.H2O + 3H2O

(II) Properties :
(a) It is a white powder.
(b) It has the property of setting to a hard mass when a paste with water is allowed to stand aside for
sometime.
(c) When it heated at 2000C, anhydrous CaSO4 is formed.
(III) Uses :

(a) In surgery for selting broken bones

(b) In making casts for toys, statues etc.
(c) In making blackboard chalks.
Anomalous behavior of Lithium :-
(a) On account of its small size exerts the greatest polarising effect out of all the alkali metals and ions,
consequently covalent character is developed.
(b) Li has the highest ionisation energy and electronegativity as compared to other alkali metals.
(c) It is not affected by air easily and does not lose it lusture even on melting.
(d) It is more harder and lighter than other alkali metal.
(e) It reacts slowly with water to liberate hydrogen.
(f) When burnt in air or oxygen, it forms only monoxide, Li2O. Howerver, the rest of the alkali metals give
peroxide or superoxides.
(g) Li2O is much less basic oxides of other alkali metals.
(h) Lithium is the only alkali metal which directly reacts with nitrogen to from Li3N.
(i) Lithium hydroxide decomposes red heat to form Li2O. Hydroxides of other alkali metals do not decompose.
2LiOH Li2O + H2O
(j) LiHCO3 is known in solution but not in solid state while the bicarbonates of other alkal metals are known
in solid state.
(k) Li2CO3 is less stable, as it decomposes on heating. Li2CO3 Li2O + CO2
(l) Li2SO4 is the only alkali metal sulphate, which does not form double salts Ex. alum.
(m) Li when heated in NH3 forms imide Li2NH while other alkali metals form amides. MNH2
(n) Lithium shows resemblance with magnesium, an element of group IIA. This resemblance is termed as diagonal
relationship.
 Shubhang Garodia
Similarities Between lithium and Magnesium :

(a) Both lithium and magnesium are harder and lighter than other elements in the respective groups.
(b) Lithium and magnesium react slowly with cold water. Their oxides and hydroxides are much less soluble
and their hydroxides decompose on heating. Both form a nitride by direct combination with nitrogen,
Li3N and Mg3N2.
(c) The oxides, LiO2 and MgO do not combine with excess oxygen to give a peroxide or a superoxide.
(d) The carbonates of lithium and magnesium decompose easily on heating to form the oxide and CO2. Solid
bicarbonates are not formed by lithium and magnesium.
(e) Both LiCl and MgCl2 are soluble in ethanol.
(f) Both LiCl and MgCl2 are deliquescent and crystallise from aqueous solution as hydrates,
LiCl.2H2O and MgCl2.8H2O.

Anomalous behaviour of Beryllium

(a) It is the hardest of all alkaline earth metal as maxiumum metallic bonding is their due to smallest size.
(b) The melting and boiling points of the beryllium are the highest.
(c) It is least reactive due to highest ionisation potential.
(d) Due to high charge density its polarising effect is highest and it forms covalent bond.
(e) It dissolves in alkalies with evolution of hydrogen
Be + 2NaOH + 2H2O Na2BeO2.2H2O + H2
Sodium beryllate
other alkaline earth metals do not react with alkalies.
(f) Oxides and hydroxides of beryllium are amphoteric in nature.
BeO + H2SO4 BeSO4 + H2O BeO + 2NaOH Na2BeO2 + H2O
Be(OH)2 + 2HCl BeCl2 + 2H2O Be(OH)2 + 2HCl Na2BeO2 + 2H2O
The hydroxide is unstable in water and covalent in nature.
(g) Like Al, its carbide (Be2C) on hydrolysis evolves methane.
(h) Due to its small size it has strong tendency to form complex.
(i) It shows diagonal relationship with Al.

Diagonal similarity between beryllium and aluminium :

In many of its properties, beryllium resembles aluminium.

Thus –
(a) The two elements have same electronegativity and their charge/ radius ratios.
(b) Both metals are fairly resistant to the action of acids due to a protective film of oxide on the surface. Both
metals are acted upon by strong alkalies to form soluble complexes, beryllates [Be(OH) 4 ]2– and
aluminates, [Al(OH)4 ]– .
Cl Cl Cl Cl
(c) The chlorides of both beryllium and aluminium Cl-Be Be-Cl Al Al
have bridged chloride structures in solid state. Cl Cl Cl Cl
 Shubhang Garodia
(d) Salts of these metals form hydrated ions, Ex. [Be(OH2)4 ]2+ and [Al (OH2)6 ]3+ in aqueous solutions. Due to
similar charge/ radius ratios of beryllium and aluminium ions have strong tendency to form complexes.
For example beryllium forms tetrahedral complexes such as BeF4 2– and [Be(C2 O4 )2 ]2– and aluminium
forms octahedral complexes like AlF6 3– and [Al(C2 O4)3]3–.
Cement :
It is a light grey, heavy fine powder, It is a homogenous mixture of silicates and aluminates of calcium, which
form more than 90% of the cement are –
(a) Tricalcium silicate - 3CaO.SiO2

(b) Dicalcium silicate (slowest setting component) - 2CaO.SiO2

(c) Tricalcium aluminate (fastest setting component) - 3CaO.Al2O3

(d) Tetracalcium alummino ferrate - 4CaO.Al2O3.Fe2O3

Raw Materials Composition of Cement

(i) Lime Stone – It provides CaO

CaO(61.5%) SiO2(22.5%)
(ii) Clay – It provides Al2O3 and silica (SiO2)
Fe2O(2.O%) Al2O(7.5%)
3
Cement 3

K2O(1.5%) 100% MgO(2.5%)

(iii) Gypsum – CaSO4. 2H2O

SO3(1.5%) Na2O(1.5%)

Setting of cement : When water is mixed to cement and the mixture is left it becomes very hard. This
property of cement is called setting.

Mortar : It is a mixture of cement, sand and water to give a proper consistency.

Concrete : A mixture of cement, Sand gravel and water is known as concrete.
Reinforced concrete cement (RCC) : When concrete is filled in beams made of iron bars, it is called RCC.
Iron imparts extra strength to the structure.

Biological role of Sodium, Potassium, Magnesium & Calcium :

(I) Role of Na+ and K+

An average human body having 0.07 kg of Na+ and 0.25 Kg of potassium. Na+ is an extracellular while K+ is an
intracellularion.
(a) Maintenance of normal hydration and osmotic pressure.
(b) Sodium and potassium are important in maintenance of normal neuromuscular irritability and excitability.
(c) In blood plasma, sodium and potassium are important in the maintenance of the proper viscosity of blood.
(d) The base in alkaline digestive juices, such as the pancreatic juice and bile, is derived from blood sodium
and potassium salts.
(e) The ratio of K+ to Na+ ions in human beings is 7 : 1. This particular ratio, called concentration gradient.
(f) Many enzyme reactions are catalysed by Na+ and K+ ions.
 Shubhang Garodia
(II) Role of Ca2+ ions

(a) Mineralisation of cartilage into bone : The main inorganic salts deposited in the organic matrix of bone
are hydroxy apatites. The replacement of the ground substance of the cartilage by hydroxyapatite, is
called as mineralisation of bone.
(b) Calcium is essential and intimately related with growth.
(c) Ionic calcium helps in the production of thromboplastin and in the conversion of prothrombin into thromin.
The removal of Ca2+ from the blood can prevent blood coagulation.
(d) Calcium is a brilliant activator for enzymes.
(e) Calcium ions help in the release of acetylcholine from the storing vesicles present in the nerve ending.
(f) Calcium is essential for the excitation of nerves.
(III) Role of Mg2+ ions

(a) Mg2+ ions serve as activators of importat enzymes including enolase, peptilase, alkaline phosphatase
RNA & DNA Polymerase.
(b) It exerts an effect on neuromuscular irritability and high level inducing depression.
(c) In the body, Mg2+ and Ca2+ act as antagonists to one other and counteract certain effects of one another.
(d) Mg2+ ions are present in chlorophyll which in turn plays an important role in photosynthesis.