JEEMAIN.

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Vidyamandir Classes s-Block

s-BlockElements
Elements&&Compounds
Compounds

s-Block Elements & Compounds

PROPERTIES OF s-BLOCK ELEMENTS Section - 1

Group - I
Introduction :
All the alkali metals have loosely held one s-electron in the outermost shell which they can readily lose to
give monovalent (M+) cation having stable noble gas configuration. Due to their tendency of loosing s-
electron easily, they have low ionization energy and high metallic character. The size of atoms and ions of
alkali metals increases down the group.

Physical Properties :
(i) Electropositive character
Alakali metals are highly electropositive in nature and electropositive character increases down the
group.
(ii) Ionization Energy
Alkali metals have low ionisation energy and it decreases down the group. In fact K and Cs are used
as cathodes in photoelectric cells.
(iii) Density
The density of alkali metals is quite low as compared to other metals. Li, Na and K are even lighter
than water. As we go down the group, the mass and volume of alkali metals increases but mass
increases by larger factor than the volume and the resulatant effect is that the density increases down
the group with an exeption that potassium beigh lighter than Sodium. So, the trend is :

Li  K  Na  Rb  Cs

(iv) Melting Point and Boiling Point :

Melthing point and boiling point of alkali metals decreases down the group. The melting points range
from lithium 181C to caesium 28.5C. These are extremely low values for metals, and contrast with
the melting point of the transition metals, most of which are above 1000C.
(v) Flame Test :
Group I elements give a varied range of colours in their flame test. Li emits crimson light, Na emits
yellow, K emits lilac and Rb and Cs emit violer light.

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Chemical Properties :
The alkali metals are highly reactive due to low ionization enthalpy and therfore they never occur in free
state. The reactivity of these metals increases down the group.
(i) Reactivity towards air
Li react with air forming oxide Li2O (and some Li2O2) and nitride Li3N. Na react with air forming
oxide (Na2O) and peroxide (Na2O2). Peroxide is formed in large amount. K, Cs and Rb forms oxide
(M2O), peroxide (M2O2) and superoxide (MO2). Superoxide is fromed in large amount. (Where M
is K, Cs, Rb).
You can note here except Li all other alkali metals are forming oxide only, whereas Li form nitride also
on buring in air. Li3N is a ruby red salt which gives LiOH and NH 3  on dissolving in water while Li
and nitrogen on simply heating. The increasing stability of peroxide or super-oxide, as size of the metal
ion increases, is due to the stabilization of large anions by larger cation through higher lattice energies.
(ii) Reactivity towards water
Group 1 metals all react vigorously with water liberating hydrogen. The reaction becomes increasingly
violent on descending the group.
1
M s   M sq  le ; H 2O  le  OH   H2
  2
1
 M   OH  
M  H 2O  H2 (M  Li, Na, K, Rb, Cs)
2
Reaction of sodium is os violent that it catches fire and is thus kept in kerosene in the laboratory.

Standard Reduction Potential

Standard Reduction potential of akali metals M is a value that represents the tendency to gain an electron

M aq  le  M s 

 
If standard Reduction potential is positive it means element desires to gain electrons and if it is negative
it means element desires to loose electron.
Further, lower the standard reduction potential, higher will be the tendency to loose electron.
Standard reduction potential (Eo) for an alkali metal represents the overall change :

M  s   M  g  Sub lim ation enthalpy

M  g   M   le  Ionization enthalpy

g 
M  M
 H 2O  hydration enthalpy
g   aq 

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Lithium has the most negative standard reduction potential than any other element in the periodic table.
Why ?
a Li  H  159 KJ /mol
ion Li  H  520.2 KJ/ mol
Lis   Li g  

 hydr H Li   520
Li   le  Li   le 
g   aq 

Total energy aborbed  aH Li  ion H Li   hydn H Li  159  520.2  520  159.2 KJ / mol.

a Na  H  107 KJ/ mol

ion Na  H  495.8 KJ/ mol
Na  s  
 Na  g  


 hydr H Na   406 K / mol

Na   le 
 Na   le
g   aq 
Total energy absorbed
 a H Na  ion H Na   hydr HNa   107  495.8  406  196.8 KJ ? mol

Hnece total energy absored in Lis   Li  aq   le  is less and also least as compared to any
other element. Therefore it has most negative standard reduction potential.

Li has most negative standard reduction ptoential (or highest tendency for Li  s   Li   aq   le  ).
It seems that reaction of Li with water should be most vigourous. But it is surprising that Li reacts less
vigorously with water than other alkali metals. The explanation lies in the kinetics (that is the rate at which
the reaction proceeds) rather than in the thermodynamics (that is total amount of energy absorbed). You
will study more about kinetics in upcoming modules.
(iii) Reducing Character
Among alkali metals, Lithium has strongest reducing character and sodium has least reducing
character and rest are almost the same.
[Lower the standard reduction potential (considering the sign also) higher is the reducing character]
(iv) Solution in liquid Ammonia
Liquid Ammonia is also a good polar solvent next the water. Ammonia gas (b.p. = - 33oC) is
condensed to give liquid ammonia. Both water and ammonia undergo self - ionization :



2H 2O  
 H3O  OH

; 

2NH3   
 NH 4  NH 2

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When small amount of sodium is added in liquid ammonia (i.e. : Dilute Solutions of Na in liquid
Ammonia), sodium metal looses an electron to give sodium cation, both of these ions are solvated by
liquid ammonia to give a blue coloured solution which is highly conducting.
 
M   x  y  NH3   M  NH3  x   e  NH3  y 
 
The blue colour of the solution is due to the ‘ammoniated electron’ which absorbs energy in the visible
region of light and thus imparts blue colour to the solution. The conducting nature is also mainly due to
solvated electron.
The solution is paramagnetic in nature and on standing slowly liberates hydrogen resulting in the
formation of amide.
1
e  NH3    NH 2 ammoniated   H 2 g 
2
Concentrated solution of Na in liquid ammonia is metallic bronze in colour and diamagnetic in nature
due to formation of metal ion clusters.

Group - II
Introduction :
All the alkaline earth metals have two s-electron in the outermost shell which if they lose, they will give
divalent (M2+) cation having stable noble gas configuration. The size of atoms and ions of alkaline earth
metals increases down the group.
Physical Properties :
(i) Electropositive character
Alkaline earth metals are highly electropositive in nature and Electropositive character increase down
the group.
(ii) Ionization Enthalpy
Second ionization enthalpy is very high than first ionization enthalpy. It is due to the fact that extracting
an electron from a positive ion bcomes difficult. Both first and second ionization energy decreases
down the group.
(iii) Density
The density of alkaline earth metals is high as compared to alkali metals. And as we go down the
group both the mass and volume of alkali metals increases but the resultant effect that is the density do
not show a regular change. As we move down the group, it first decrease (upto Ca) and then increases.
So, the trend is : Ba > Sr > Be > Mg > Ca

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(iv) Melting point and Boiling Point

Melthing point and boiling point of alkaline earth metals is higher than the alkali metals. As we move
down the group,
Melthing point and Boiling point don’t show a regular change.

Chemical propeties :
(i) Reactivity towards air
Alakaline earth metals react with air forming an oxide and a nitride.

air
M s   MO  M 3 N 2  where M  Be, Mg, Ca, Sr, Ba 

The case of above reaction depends on electropositive character and hence increases down the
group. In the case of Mg, it burns with a dazzling white light which is used to provide light in flash
photography using light bulks.

(ii) Reactivity towards water


 M 2aq
OHC : Ms    2e  RHC : 2H 2O    2e 
 H 2  2OH 
 
Now, to check the case of reaction for different metals, we have to check the Oxidation Half Cell
reaction and this reaction refers to the reducing ability (Standard Reduction Potential) of the metal.
And, Since for alkaline earth metals reducing character increases down the group, the case of reaction
of alkaline earth metals with water increases down the group.
Experimentally, it is see that Be reacts only with steam and Magnesium can react with both hot water
and steam and Ca. Sr, Ba ract even with cold water rapidly.

(iii) Solution in liquid ammonia

In liquid ammonia, group II metals form bright blue dilute solutions containing solvated electrons and
metal hexaammoniates. The metal hexammoniates form ammides on heating which further form nitrides
and NH3 concentrated solution are bronze coloured.

(iv) Reducing character

Less is the standard reduction potential more will be reducing character and since the standard reduction
potential decreases down the group, the reducing character of alkaline earth metals increases down
the group.

NOW ATTEMPT IN-CHAPTER EXERCISE-A BEFORE PROCEEDING AHEAD IN THIS EBOOK

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IMPORTANT CONCEPTS Section - 2

In this section we will discuss some important concepts related to s-block elemets.

Lattice enthalpy
It is the energy required to separate one mole of a solid ionic compound completely in gaseous ions.
Lattice Enthalpy of breaking solid NaCl in gaseous Na+ and Cl– is + 788 KJ/mol.
In a system of + ve and -ve ion the energy required to separate the two ions is inversely proportion to the
distance between the centres of the two ions.

1
 LE 
r  r 

(r+ is the radius of +ve ion and r– is the radius of -ve ion.)

For a given anion and different cations as we move down the group the size of cation increase and conse-
quently the interionic distance between cation and anion increases. Hence, the lattice enthalpy would de-
crease dwon the group.
Also you may not that as the charge of ion increases the Lattice Enthalpy increases. Therefore, Lattice
Enthalpy of group II metal salt is more as compared to corresponding group I metal salts.

Hydration Enthalpy
Hydration Enthalpy refets to the energy released when one mole of gaseous ion is dissolved (or hydrated) in
water. Strictly speacking, value of Hydration Enthalpy is equal to the energy absorbed when one mole of
gaseous ion is dissolved in water. For example, if energy is released on dissolving some ion in water, its
hydration enthalpy will be negative quantity and its magnitude will be equal to the magnitude of energy
released.
When an ion is dissolved in water it attracts water molecules. Thus, a number of water molecules surround
it and a cluster is formed. The number of water molecules surrounded by a cation depend on the ability of
the cation to polarise the water molecule. More the polarizing power of cation, the large the number of
water molecules will surround it. Among alkalie metals cations, Li+ has highest polarising power, hence, a
large number of water molecules will surround it. The numbere is so large that the hydrated Li+ becomes
heavier than hydrated Na+. The hydrated Na+ being lighter moves easily in water as compared to hydrated
Li+, hence Na+ is more conduction than the solution of Li+.

Also, note that more that polarising power of cation, more it will be hydrated and more energy will be
released. And therefore, hydration enthalpy will be more negative. Thus down the group magnitude of
hydration enthalpies of metal ions decreases.

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Thermal Stability
Hereby we will discuss the thermal stability of group I and group II metal oxosalts. i.e. carbonates, nitrates,
sulphates etc. For simplicity let us consider metal carbonates and discuss their thermal stability.

Example :
The effect of Heat on the metal carbonates :
All the carbonates undergo thermal decomposition to give the metal oxide and carbon dioxide gas. Thermal
decompostion is the term given to splitting up a compound by heating it.

XCO3  s   XO  s   CO 2  g 

As you go down the group, the carbonates have to be heated more strongly before they will decompose.
i.e. Thermal stability of metal carbonates increase down the group. Let us try to understand the explanation
of this fact.

Explanining the trend the terms of the polarising ability of the positive ion :
A small cation has a lot of charge packed into a small volume of space. It has a high charge density and will
have a marked distorting effect on any negative ions which happen to be near it.
A bigger cation has the same charge spread over a larger volume of space. Its charge denisty will be lower,
and it will cause less distortion to nearby negative ions.

The structure of the carbonate ion :

If you worked out the structure of a carbonate ion, you would probably come up with :

This show two single carbon-oxygen bonds and one double one, with two of the oxygens each carrying a
negative charge. Unfortunately, in real carbonate ions all the bonds are identical, and the charges are spread
out over the whole ion - although concentrated on the oxygen atoms. We say that the charges are delocalised.
The next diagram shows the delocalised electrons. The shading is intended to show that there is a greater
chance of finding them around the oxygen atoms than near the carbon.

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Polarisin the carbonate ion :

Now imagine what happens when this ion is placed next to a positive ion. The positive ion attracts the
delocalised electrons in the carbonate ion towards itself. The carbonate ion becomes polarised.

If this is heated, the carbond dioxide breaks free to leave the metal oxide.
How much you need to heat the carbonate before that happens depends on how polarised the ion was. If it
is highly polarised, you need less heat than if it is only slightly polarised.
The smaller the positive ion is, the higher the charge density, and the greater effect it will have on the
carbonate ion. As the positive ions get bigger as you go down the Group, they have less effect on the
carbonate ions near them. To compenstate for that, you have to heat the compound more in order to
persuade the carbond dioxide to break free and leave the metal oxide. In other words, as you go down the
Group, the carbonates become more themally stable.
In case of Nitrates, Group II metal nitrates decompose to give metal oxide, nitrogen dioxide and oxygen gas
while Group I metal nitrates except lithium) decompose to give metal nitrite and oxygen gas. Lithium nitrate
behaves in a way similer to group II metals giving. LiO, NO2 and O2 on decopostion.


Ca  NO3 2  CaO  NO 2  O 2

Li(NO3 ) 2  Li 2O  NO 2  O 2
500C
NaNO3 
 NaNO 2  O 2
800C
NaNO3 
 Na 2O  N 2  O 2

NaNO3 on very strong heating gives Na2O, N2 and O2. Metals sulphates decompose to give metal oxide,
SO2 and O2. Thermal stability of all these oxosalts increases down the group and the explanation lies in a
way similar to as given for metal carbonates.

NOW ATTEMPT IN-CHAPTER EXERCISE-B BEFORE PROCEEDING AHEAD IN THIS EBOOK

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IMPORTANT COMPOUNDS OF s-BLOCK Section - 3

1. Oxides, Peroxide, Superoxides

Gruop - I
Alkali metals generally burn in air to form oxides, peroxides and superoxides (Lithium forms some nitride
also). Lithium forms Li2O (and some Li2O2 ) and Li3N. Sodium form Na2O2 (and some Na2O), and rest of
the alkali metals form superoxides (major). The increasing stability of oxides, peroxides and superoxides, as
the size of the metal ion increases, is due to the stabilisation of large anions by larger cations through lattice
energy effects. The oxides and peroxides are colourless but the superoxides are orange or yellow. Oxides
and peroxides ae diamagnetic while superoxides are paramagnetic. Alkali metal oxides, peroxides and
superoxides dissove readily in water to give hydroxides along with a lot of heat. Peroxides and superoxides
are good oxidants and generally used in bleaching.

1. Sodium Oxide [Na2O]

Preparation :
(i) Controlled oxidation of sodium in air gives Na2O
(ii) Industrially, Na2O is prepared by heating sodium nitrate of nitrite with sodium.

NaNO3  Na 
 Na 2O  N2 ; 2NaNO2  6Na 
 4Na 2O  N2

(iii) In laboratory pure Na2O is formed by heating mixture of sodium azide and sodium Nitrite.

3NaN3  NaNO2  2Na 2O  5N 2
(pure)

Porperties :
Na2O is a white ionic solid and its aqueous solution is a strong base.

Na 2O  H 2O 
 2NaOH

2. Sodium Peroxide [Na2O2]

The trade name of sodium peroxide is oxone.
Preparation :
Burning Sodium in air mainly forms sodium peroxide.

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Properties :
(i) Na2O2 is colourless in pure form and stable at room temperature.
(ii) Na2O2 is diamagnetic. All peroxides are regarded as salts of dibasic acid H2O2.

Na 2O2  2H2O 
 2NaOH  H2O 2

(iii) Cold dilute acids always produce H2O2.

Cold
Na 2 O2  HCl  NaCl  H 2O2

The reaction with CO2 is used to purify air in submarinas but KO2 is even better for this purpose.

(iv) It is strong oxidising agent. It oxidises Al to Al2O3 , Cr 3 to CrO42  and SO 2 to SO 42  .

Na 2O2  Al 
 Al2O3  Na 2O
Na 2O 2  CO 2 
 Na 2SO4

The reaction with CO2 is used to purify air in submarines but KO2 is even better for this purpose.

Uses :
It is a powerful oxidant and used for bleaching wood pulp, paper and fabrics.

3. Potassium Superoxide [KO2]

Preparation :
It is prepared by burning potassium in excess of oxygen.

K  O 2  excess   KO 2

Porperties :
(i) It is paramagnetic and orange coloured solid.
(ii) Potassium superoxide is stronger oxidizing agent and gives both H2O2 and O2 with either water
or acids.

2KO2  2H 2O 
 2KOH  H 2O2  O 2 ; 2KO 2  2HCl 
 2KCl  H 2O 2  O 2

(iii) It readily reacts with CO and CO2 producing Oxygen.

2KO 2  CO 
 K 2CO3  O 2 ; 2KO 2  CO 2  K 2CO3  3 / 2O2

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User :
Being oxidant, KO2 produces oxygen and simulataneoulsy it absorbs CO2 and becuase of its above
life supporting properties it is used in space capsules, submarines and breathing masks.

Gruop - II
Alkaline earth metals burn in air to give normal metal nitrides. Generally, there are two methods of forming
oxides, one is buring in air and other decomposing their oxosalts.
Metal oxides and hydroxides of group II are more stable than of group I, that is why oxosalts (carbonates,
sulphates, nitrates,....) of group II are less stable to heat and decompose to give corresponding oxides.

Note : The stability of metal oxides decreases down the group. That is why, down the group metal oxides become
more reactive. Down the group, alkaline earth metal oxide’s basic character increases. BeO is amphoteric
in nature.

1. Magnesium Oxide
Preparation
(i) Magnesium burns in air with a dazzaling flame and forms magnesium oxide.

2Mg  O 2 
 2MgO

(ii) Decomposing the oxosalt MgCO3 also gives magnesium oxide.


MgCO3 
 MgO  CO2

Properties :
(i) It is light infusible (fusibility refers to conversion in liquid form) white powder. It fuses at 2800oC.
(ii) It is reduced by carbon at very high temperature.
2000C
MgO  C 
 Mg  CO

Uses :
(i) Buring of magnesium ribbon is used to initiate the thermite reaction.The thermite reaction goes
as :

Fe 2O3  2Al   2Fe  Al2O3

(ii) MgO being very less reactive and having high m.p., is used as a refractory material. Good
conductivity of MgO towards heat and bad conductivtiy towards electricity also adds to its
usefulness as refractory material.

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2. Calcium Oxide [CaO] (Quick lime)

Prepartion :
It is prepared by decompostion of CaCO3 in lime kilns (900oC).


CaCO3  CaO  CO2

Properties
(i) Calcium oxide is a white amorphous solid, having high melting point of about 2870K.
(ii) Being a basic oxide, it combines with acidic oxides at high temperaure.

CaO  SiO 2  CaSiO3 ;  2Ca 3  PO 4 2

6CaO  P4O10 

300C
(iii) CaO  3C 
 CaC2  CO 

Uses :
(i) It is used in metallurgy to remove phosphates and silicates as slag.
(ii) By mixing with SiO2 and alumina or clay it is used to make cement.
(iii) It is used for softening of water.

2. Hydroxides

Gruop - I
Hydroxides of alkali metals are strong base. Their basic character increases down the group. They
dissolve readily in water giving much heat due to intense hydration. If we go down the group solubility
of alkali metal hydroxides increases.

1. Sodium hydroxide [NaOH] (Caustic Soda)

Preparation :
(i) Heating 10%Na2CO3 Sol. with mild of lime (Lime - Caustic Soda process) :



Na 2CO3  Ca(OH)2 
 2NaOH  CaCO3

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(ii) Electrolysis of brine :

Commercially, NaOH and Cl2 are manufactured simultaneously by the electrolysis of brine (NaCl
solution).
Reactions occuring at anode and cathode are :



NaCl  
 Na  Cl


At Anode : 2Cl 
 Cl2  2e

At Cathode : Na  le   Na

2Na  2H2O  2NaOH  H2

Another reaction may occur at the anode to a small extent.

4OH  
 O 2  2H 2O  4e 

Electrolysis is carried in either of two types of cell today, diaphragm and mercury cathode cells.
NaOH made in this way always contanins some amount of NaCl.

(a) Diaphragm cell / Nelson : A porous diaphragm of asbestos is used to keep the H 2 and Cl2
gases separated otherwise they react in an explosive chain reaction in dylight. Diaphragm also
seperates the carbon anod and cathod.

(b) Mercury Cathod Cell / Castner - Kellner Cell : The anode is made up of carbon and the
cathode of mercury. The reaction occuring at Hg cathode are thus :

Cathod : Na   e  
 Na

Na  Hg  Na  Hg amalg am 

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The amalgam is pumped in a different comartment called denuder where water trickles over.

1
Na  Hg  H 2O 
 NaOH  H 2  Hg  recycled back to electrolysis tan k 
2

Properties :
(i) It is a white crystalline solid, deliquiscent (absorbs moisture from air), soluble in water, good conductor
and gives silky touch like soap.
(ii) It reacts with acid and acidic oxides to form salts

NaOH  HCl 
 NaCl  H 2O
2NaOH  SiO 2 
 Na 2SiO3  H 2O
2NaOH  CO 2 
 Na 2CO3  H 2O

(iii) It reacts with amphoteric metals such as Sn, Al, Zn and amphoteric metal oxides.

Zn  2NaOH   Na 2 ZnO 2  H 2 Sodium Zincate 

3
Al  NaOH  H 2O 
 NaAlO 2  H 2  Sodium Aluminate 
2
Sn  2NaOH 
 Na 2SnO 2  H 2  Sodium Stannate 

SnO  2NaOH 
 Na 2SnO2  H 2O
SnO2  2NaOH 
 Na 2SnO3  H 2O
Al2O3  2NaOH 
 2NaAlO2  H 2O

(iv) It reacts with ammonium salts to give ammonia, whcih serves as a test for ammonia.

NH 4Cl  NaOH 
 NaCl  NH3   H 2 O

The above reaction is double displacement reaction only. NH4OH must be formed as second product.
But NH4OH is never formed as a product. It decomposes to give NH3 + H2O.
(v) Disproportion of some non-metals in NaOH : (Learn these reactions)

X2  NaOH 
 NaX  NaOX  H 2O  X  Cl, Br, I
 Cold & Conc.

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X2  NaOH 
 NaX  NaOX3  H 2O
 hot and Conc.

P4  NaOH  
 NaH 2 PO 2  PH 3
 phosphine 

[This reaction is used in the preparation of phosphine gas.]

S  NaOH 
 Na 2S  Na 2S2O3  H2O

(All above reactions can be balanced as Redox Reactions)

(vi) Reactio with salts of transition metals :

 Fe  OH 3  3NaCl
FeCl3  3NaOH 

When hydroxides are unstable, the oxides are precipitated.

 HgO     H 2O  2NaCl
HgCl2  2NaOH 

Hg  OH  2

 Ag 2O     H 2O  2NaNO3
2AgNO3  2NaOH 
  
AgOH

2. Potassium Hydroxide [KOH] (Caustic Potash)

Preparation :
It can be prepared by electrolysis of KCl solution similar to electrolysis NaCl solution (brine)

At Cathode : 2K   2H 2O  2e  
 2KOH  H 2 

At anode : 2Cl  2e 

 Cl2 

Properties :
(i) The properties are similar to those of NaOH but being expensive is less used.

(ii) KOH is more soluble in alcohol than NaOH and produces OC2H5 ions.

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 


C2 H5OH  O H 
 OC2 H5  H 2O

This accounts for use of elcoholic KOH in OrganicChemistry which you will study in modules of organic.
Uses :
(i) It is used in organic chemistry
(ii) KOH is used as an absorbent for CO2, for which is perferred over NaOH, since after absorption
of CO2, the KHCO3 formed is solube whereas NaHCO3 being sparingly soluble separates out,
and chokes the absorption bulbs.

Gruop - 2
Alkaline earth metal oxides react with water to form sparingly soluble metal hydroxides.
Alkaline earth metal hydrooxides are less basic and more stable than alkali metal hydroxides. The solubility,
thermal stability and basic character of the hydroxides increases down the group. Beryllium hydroxide is
amphoteric in nature as it reacts with both acids and bases.

Be  OH 2  2NaOH 
 Na 2  Be  OH 4 
Sodium beryllate

Be  OH 2  2HCl  2H 2O 
  Be  OH 4  Cl2

Rest of the alkaline earth metal hydroxide are basic in nature.

1. M agn esi u m H y d r o x id e [ M g(OH ) 2]

Preparation :
Mg(OH)2 is pepared by dissolving magnesium oxide in water.

 Mg  OH 2
MgO  H 2O 

Properties :
(i) It is a white powder extermely insoluble in water.
(ii) It is weakly basic and suspension Mg(OH)2 in water [milk of Magnesia] is used as an antacid.

2. Calcium Hydroxide [CaOH)2]

Preparation :
Calcium hydroxide is prepared by adding water to quick lime (CaO).

 Ca  OH  2
CaO  H 2O 

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Properties
(i) It is white powder, sparingly soluble in water.
(ii) The aqueous sturated solution of Ca(OH)2. is known as lime water and the suspension of slaked
lime (solid Ca(OH)2) in water is known as milk of lime.
(iii) Bleaching powder is formed when chloring gas is passed through lime water.

3Ca(OH)2  2Cl2 
 Ca(OCl)2 . Ca(OH)2 . CaCl2 . 2H2O

Uses :
(i) It is used in white wash.
(ii) It is used in the solvay’s process.
(iii) It is used to remove temperaray hardness of water :

Ca(HCO3 ) 2  Ca  OH 2 
 2CaCO 3   2H 2O

Mg(HCO3 )2  Ca(OH)2 
 2CaCO3   Mg(OH)2   2H2O

(iv) When carbon dioxide is bubbled through lime water it turns milky and when excess carbon dioxide is
passed, the precipitate (CaCO3) dissolves forming Ca(HCO3)2.

CO
2  CaCO () 
2  Ca(HCO ) CO
Ca(OH)2(aq)  3 3 2(aq)
white precipitate Excess
soluble

Similaraly, Ba(OH)2 (Baryta water) responds, when CO2 is bubbled through it.

CO
2  BaCO () 
2  Ba(HCO ) CO
Ba(OH)2  3 3 2(aq)
milky so ln. Excess
milkiness disappears

In this reaction Ba(OH)2 is even more sensitive ot CO2 but Ba(OH)2 being more expensive,
Ca(OH)2, is preffered to detect CO2 in laboratory.

Note : The caustic alkalis (NaOH and KOH) are the strongest bases known in aqueous solution. Soda lime is a
mixture of NaOH and Ca(OH)2 and is made from quick lime (CaO) and aqueous NaOH. Soda lime is
much easier to handle than NaOH.

NOW ATTEMPT IN-CHAPTER EXERCISE-C BEFORE PROCEEDING AHEAD IN THIS EBOOK

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3. Halides
Gruop - I
Alakali metal halides like NaCl and KCl are generally found in sea water. Among Sodium halides,
stability, melthing point and boiling point decreases down the group. Solubility for alkali metal halides
doesnot show a regular trend.
1. NaCl
Preparation :
It occurs in sea water as deposits in mines as rock salt. It is extracted from sea water simply by
solar evaporation. The crystals of NaCl are washed to remove more soluble MgCl2.
Properties :
(i) It is colourless crystalline (in pure form). it is NOT hydroscopic but pressence of MgCl2
makes it deliquescent due to the high polarizing power of MgCl2 (that attracts H2O molecules)
(ii) It is used to produce industrially important compounds like Na2CO3 and Cl2.
(a) Na2CO3 is produced by Solvay’s Process. (discussed later)
(b) Cl2 is produced by :
(i) Electrolysis (NaOH is also a product) of brine.
(ii) Leblanc Process : NaCl  conc. H 2SO4 
 NaHSO4  HCl
heat
NaHSO4  NaCl 
 Na 2SO4  HCl

 Cl2  Mn 2 
HCl  MnO2 

Uses :
(i) Eating salt in food.
(ii) It is used to lower the melting point of ice. Kulfi seller uses NaCl to sustain Kulfi for long time
in summer.
KCl is also extracted from sea water. Its properties are almost similar to NaCl. It has an important
use in producing fertilizers.

Gruop - II
Alkaline earth metal halides are generally made by heating matals with halogen or by action of halogen
acid on metal or metal or metal carbonates. Beryllium halides are covalent while rest of the alkaline
earth metals form ionic halides.
BeCl2 has a chain structure as shown in the figure.

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Here Cl bonded to one Be uses a lone pair of e–s to form a coordinate bond to another Be atom.
Dehydration of hydrated chlorides, bromides and iodies of Ca, Sr, Ba can be achieved on heating but the
hydrated halides of Be and Mg suffer hydrolysis on heating. In aqueous solution BeCl2 , Be exists as
2
 Be  H 2O   or  Be  H 2O 4  Cl2 .
 4

heat
 Be  H 2O   Cl2   Be  OH  2  2HCl
 4

1. Magnesium Chloride [MgCl2 . 6H2O]

Preparation
(i) It is prepared by passing dry HCl over Magnesium.

Mg  2HCl 
 MgCl2  H 2 

(ii) Dow’s process :

The extraction of Magnesium from sea water depends on the fact that Mg(OH)2 is very much less
soluble than Ca(OH)2.

Ca(OH)2  MgCl2 
 Mg (OH)2   CaCl2
(in sea water )
acidified
Mg(OH)3  MgCl2(aq) ( MgCl2 . 6H 2O)
with HCl
dry HCl
MgCl 2 . 6H 2 O  MgCl2  6H 2 O
Cl 2 /SOCl 2

Note : If MgCl2 . 6H2O is heated, then the products are MgO and HCl. So water of crystallisation are removed
by passing dry HCl/Cl2 or SOCl2.

Properties
(i) It is colourless solid, highly soluble in water.
(i) When heated, it decomposes

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
MgCl2 . 6H2O 
 Mg(OH) Cl  HCl  5H2O

Mg(OH) Cl 
 MgO  HCl

Uses :
It is used in electrolytic method for extracting Magnesium.
2. Calcium Chlorid [CaCl2 . 6H2O]
Preparation
It is formed as Solvay’s Process’ by-product.
Properties
(i) It reduces the freezing pt. of water.
(ii) Anhydrous salt is an excellent drying agent. It forms addition compounds with NH3 and CH3OH
(CaCl2 . 8NH2 and CaCl2 . 4CH2OH) and hence cannot be used to dry them.
Uses :
It is sprinkled on roads in hilly areas to remove snow.

4. Carbonates :
Gruop - I
Alakli metal carbonates are basic salts. The stability of alkali metal carbonates towards heat increases
down the gourp. Thye quite stable and melt before they eventually decompose into oxides (at above
1000oC).
1. Sodium Carbonate (Na2CO3) :
Washing soda is Na2CO3 . 10H2O
Soda ash in Na2CO3
Preparation :
Solvays’s Process (Ammonia-Soda Process)
The process is much more complicated than the overall equation and since the reactions involved
are reversible only 75% of the NaCl is converted.
The purified brine (NaCl) solution is first saturated with ammonia and then carbonated with CO2
forming NaHCO3

 HCO3  H 
CO2  H2O  ; NH3  H 
 NH 4
NH 4  HCO3  NaCl 
 NaHCO3  NH 4Cl

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The NaHCO3 formed is insoluble in the brine solution becasue of the common ion effect and so can
be filtered off. This is then heated to form anhydrous Na 2CO3 .

150C
2NaHCO3  Na 2CO3  H2O  CO 2
NH4Cl produced above reacts with lime water forming a very good dehydrating agent (CaCl2) and
evolving NH3 gas.

2NH 4Cl  Ca  OH 2 
 CaCl 2  2NH 3  2H 2O

Lime water used above was prepared instantaneously by heating CaCO3 and passing water through
quick lime.

CaCO3  CaO  CO 2 ;  Ca  OH  2
CaO  H 2O 

The materials consumed are NaCl and CaCO3 and the useful product is Na2CO3. CaCl2 which is a
by prodcut is little used and the rest is wasted. CO2 and NH3 formed are used again to continue the
process.
The whole process can be diagrammatically shown as :

Properties :
(i) On passing CO2 through aqueons solution of Na2CO3, NaHCO3 is formed.

Na 2CO3  CO2  H2O 

 2NaHCO3 

(ii) When aqueous solution of sodium carbonate containing sulphur is treated with sulphur dioxide,
sodium thiosulphate is formed.

Na 2CO3  SO2  H2O 

 Na 2SO3  CO2 

Na 2SO3 + S 
 Na 2S2O3

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(iii) On heating washing soda (Na 2CO3 10 H 2O) it forms monohydrate called heavy ash or crystal
carbonate.


Na 2CO3 . 10H 2O 
 Na 2CO3 . H 2O  9H 2O

Uses :
(i) Used in water softening and cleaning.
(ii) Used in paper, paint and textile industries
Potassium carbonate (Pearl Ash) is not much important. Its properties resemble closely with
Na2CO3. It is prepared by carbonation of caustic potash.

KOH  CO2 
 K 2CO3  H2O

Note : K2CO3 cannot be prepared by Solvay’s Process because KHCO3 formed in the reaction is highly soluble
and hence cannot be separated form (NH4Cl + KHCO3) mixture easily.

Gruop - II
Carbonates are basic salts. BeCO3 is covalent while other carbonates are ionic. On heating, alkaline earth
metal carbonates forms an oxide and carbon dioxide is evolved.
1. Calcium Carbonate [CaCO3] :
Preparation :
It is mainly extracted from its ores. It can be prepared by these methods also :
(i) It can be prepared by passing limited CO2 through lime water.

Ca(OH)2  CO 2 
 CaCO3  H 2O
(ii) It can be obtained by adding sodium carbonate solution to CaCl2.

CaCl2  Na 2CO3 
 CaCO3  2NaCl

Properties :
(i) It is white powder, almost insoluble in water.
1200 K
(ii) On heating at 1200 K, CO2 is evolved CaCO3   CaO  CO2 (  )

(iii) CaCO3  HCl 

 CaCl2  H2O  CO2 (  )



H2CO3

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Uses :
(i) It used in manufacture of cement and quick lime.
(ii) Used in Solvay’s process in the manufacture of Ca(OH)2.
(iii) It is used in toothpaste.

5. Bicarbonates
Gruop - I
Alkali metal are strongly basic and form solid bicarbonates (execpt lithium) while no other melal forms
solid bicarbonates LiHCO3 can exist only in solution. Alkali metal hydrogen carbonates are soulble in
water. On heating, they decompose to give carbonates.

The hydrongen bonding in bicarbonates enhances the stability. In NaHCO3, the HCO3 ions are
linked to form infinite chain while in KHCO3, a dimeric anion is formed.

As the electropositive character increase down the group, the stability of hydrogen carbonates also
increases down the group.
1. Sodium Bicarbonate [NaHCO3] (Baking Soda)
Preparation :
It is an intermediate product of Solvay’s process.

NaCl  NH 4  HCO3 
 NaHCO3  NH 4Cl

Properties :
On heating, it decomposes to give Na2CO3, evolving CO2 which is used to detect bicarbonates.


NaHCO3 
 Na2CO3  H2O(  )  CO2 (  )

Uses :
(i) It used in fire extinguisher.
(ii) It is used as baking powder in manufacturing of cake making it fluffy. Backing powder contains
NaHCO3 and Ca(H2PO4)2

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50  100C
2NaHCO3  Na 2CO3  H 2O  CO 2 
Ca  H 2 PO 4  2  NaHCO3 
 CO 2 

The CO2 produced makes cake or bread rise.

6. Sulphates
Gruop - I
Sodium sulphate (Na2SO4 . 10H2O) is named as Glauber’s salt. It is produced as a by-product while
manufacturing HCl.


NaCl  H 2SO 4 
 NaHSO4  HCl  ; NaHSO 4  NaCl 
 Na 2SO 4  HCl 
 conc. salt cake 
Fomation of hydrated salts of Na2SO4 :

T  32C T  12C
Na 2SO4 
 Na 2SO 4 .10H 2O ; Na 2SO 4 
 Na 2SO4 . 7H 2O

It is used in paper industry.

Gruop - II
Sulphates of the alkaline earth metals are all white solids & stable to heat. Be and Mg sulphate are
highly soluble while other alkaline earth metal sulphates are very less soluble. This is due to very high
hydration enthalpy of Be2+ and Mg2+. Thermal stability of alkaline earth metal sulphates increases
down the group while solubility decreases.
1. Magnesium Sulhate [MgSO4]
It is a colourless soluble salt whcih decomposes to MgO, SO2 and O2 on heating.
MgSO4 . 7H2O is called Epsom salt while MgSO4.H2O is called Kieserite.

2. Calcium Sulphate [CaSO4]

Preparation :
It is prepared in lab by adding dil. H2SO4 to the solution of calcium salt.

CaCl2  H 2SO 4  dil  

 CaSO 4  2HCl

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Properties :

In nature it occurs as Gypsum CaSO 4 . 2H 2O.

150C 1 200C 1100C 1

CaSO 4 .2H 2O  CaSO4 . .H 2O   CaSO 4   CaO  SO 2  O 2
2 2
gypsum Plaster of paris anhydrite

When powdered plaster of paris CaSO4 . 1/2 H2O is mixed with the correct amout of water it sets
into a solid mass of CaSO4 . 2H2O (gysum). The process is used in plastering walls and plasters for
fractures.

ALUMS
The general representation of alums is [(M1 ) 2 SO 4 ] [(M III ) 2 (SO 4 )3 ]. 24 H 2O where M I represents the
metal with +1 Oxidation state and MIII represents the metal with +3 oxidation state. These alums are double
salts, which dissolve in water to give [M I  H 2O 6 ] , [M III  H 2O 6 ]3 and SO 42 ions and there-

fore., the alums are also represented as [M I  H 2O 6 ], [M III  H 2O 6 ] (SO 4 ) 2 . They form octahedral
crystals.

If MI is K  and M III is Al3 , Then the alum is potash alum  K2SO4 , Al2 (SO4 3 . 24H2O

     III
or [K(H 2O)6 ] [Al(H 2O)6 ] (SO 4 ) 2 M1 can be K , NH 4 , Rb , Cs , TL and M can be

Fe3 , Al3 , CO3 , Ga 3 , Mn 3 .  NH 4  2 SO 4 .Fe2 SO 4 3 .24H 2O is called ferric ammonium alum

while K 2SO4 .Cr2  SO 4 3 .24H 2O is chrome alum. Potash alum on heating dissolves in its own water of

crystallisation and on further heating forms K 2SO 4 .Al2  SO4 3 called burnt alum.

Micro Cosmic Salt [Na(NH4)HPO4]

Microcosmic salt is a white crystallion solid and is prepared NH4Cl and Na2HPO4 as follows :

 Na  NH 4  HPO4   NaCl
NH 4Cl  Na 2HPO 4 

On decomposition the following is obtained :


Na  NH 4  HPO 4 
 NH 4   NaH 2 PO 4

 NaPO3  Calgon   H 2O
NaH 2 PO 4 

Microcosmic salt can be used as a substitue for borax in the Bead Test

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Similar salt are also formed by Magnesium :

 Mg  NH 4  PO 4
MgSO 4  NH 4 OH  Na 2HPO4 
 
Mg  NH 4  PO 4 
 NH3  MgHPO 4  Mg 2P2O7  H 2O

IMPORTANT THINGS TO REMEMBER Section - 4

Anomalous properties of Lithium :

The anomalour behaviour of lithium is due to the : (i) exceptionally small size of its atoms and ion, and (ii)
high polarizing power (i.e., charge/radius ratio). As a result, there is increased covalent character of lithium
compounds which is resposible for their solubility in organic solvents. Further, lithium showns diagonal
relationship to magnesium which has been discussed subsequently.
Points of Difference between Lithium and other Alkali Metals :
(i) Lithium is much harder. Its M.P. and B.P. are higher than the other alkali metals.
(ii) Lithium is least reactive but the strongest reducing agent among all the alkali metals. On combustion in
air it forms mainly monoxide, Li2O and the nitride, Li3N while other metals form only oxide.
(iii) LiCl is deliquescent and crystallizes as a hydrate, LiCl.2H2O whereas other alkali metal chlorides do
not form hydrates.
(iv) Lithium hydrogencarbonates being unstable is not obtained in the solid form while all other elements
form solid hydrogencarbonates.
(v) Li2CO3, LiNO3,LiOH all form oxide on gentle heating, throgh the analogus compounds of the rest of
the group are stable.

4LiNO3 
 2Li2O  4NO 2  O 2
2NaNO3 
 2NaNO2  O2

(vi) Li2CO3, LiF and Li2O are comparatively much less soluble in water than the corresponding com-
pounds of other alkali metals.
(vii) Lithium is much heavily hydrated than those of the rest of the group.

Points of Similarities between Lithium and Magnesium :

The similarity between lithium and magnesium is particularly striking and arises because of their similar sizes
: atomic radii, Li = 152 pm, Mg = 160 pm; ionic radii : Li+ = 76 pm, Mg2+ = 72 pm. The main points of
similarity are:

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(i) Both lithium and magnesium are harder and lighter than other element in the respective group.
(ii) Lithium and magnesium react slowly with water. Their oxides and hydroxides are much less soluble
and their hydroxides decompose on heating. Both form a nitride, Li3N and Mg3N2, by direct combi-
nation with nitrogen.
(iii) The oxides Li2O and MgO do not combine with excess oxygen to give any superoxide.
(vi) The crbonates of lithium and magnesium decompose easily on heating to form the oxides and CO2.
Solid hydrogencarbonates are not formed by lithium and magnesium.
(v) Both LiCl and MgCl2 are soluble in ethanol.
(vi) Both LiCl and MgCl2 are deliquescent and crystallize from aqueous solution as hydrates. LiCl.2H2O
and MgCl2.8H2O

Anomalour Behaviour of Beryll ium :

Beryllium, the first member of the Group 2 metals, shows anomalous behaviour as compared to magnesium
and rest of the members. Further, it shows diagonal relationship to aluminium which is discussed subse-
quently.
(i) Beryllium has exceptionally small atomic and ionic sizes and thus does not compare well with other
members of the group. Because of high ionization enthalpy and small size it forms compounds which
are largely covalent and get easily hydrolysed.
(ii) The oxide and hydroxide of beryllium, unlike the hydroxides of other elements in the group, are
amphoteric in nature.

Diagonal Relationship between Beryllium and Aluminum :

The ionic radius of Be2+ is estimated to be 31 pm; the charge / radius ration is nearly the same as that of the
Al3+ ion. Hence beryllium resmebles aluminium in some ways. Some of the similarities are :
(i) Like aluminium, beryllium is not readily attacked by acids beacuse of the presence of an oxide film on
the surface of the metal.
(ii) Beryllium hydroxide dissolves in excess of alkali to give a beryllate ion, [Be(OH)4 ]2  just as alu-

minium hydroxide give aluminate ion, [Al(OH) 4 ] .

(iii) BeCl2 and AlCl3 exist in form to chain. BeCl2 form polymeric chain (chain with a large no. of BeCl2
molecules) and AlCl3 forms dimeric chain (chain with two AlCl3 molecules).
(iv) Beryllium and aluminium ions have strong tendency to form complexes, BeF42  , AlF63 .

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IN-CHAPTER EXERCISE - D

1. NaHCO3 and NaOH can not exist together in solution. Why ?

2. The hydroxide and carbonates of Na and K are easily soluble in water while the corresponding salts
of Mg and Ca are sparingly soluble in water. Explain.
3. Solvay Process is used to manufacture sodium carbonate but it is not extended to the manufacture of
potassium carbonate. Why ?
4. Why are MgO and BeO used for the lining of steel making furnance.
5. On the treatment with cold water, an element (A) reacted quietly, liberating a colourless, odourless
gas (B) and a compound (C). Gas (B) further reacts with element (A) to yield a solid product (D)
which reacted with water to give a basic solution (E). (E) is found to be same as (C). When carbon
dioxide was bubbled through solution (C) initially a white precipitate (F) is formed, but this redissolved
forming solution (G) when more CO2 was added. Precipitate (F) was heated at 1000C , a white
compound (H) was formed which when heated with carbon at 1000C , gave a solid (I) of some
commercial importance. Name the substances (A) to (I).

NOW ATTEMPT IN-CHAPTER EXERCISE-D FOR REMAINING QUESTIONS

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HYDROGEN Section - 5

Postion of Hydrogen in the Periodic Table

Hydrogen has electronic configuration 1s1. On one hand, its electronic configuration is similar to the outer
electronic configuration (ns1) of alkali metals, which belong to the first group of the periodic table. On the
other hand, like halogens (with ns2 np5 configuration belonging to the seventeenth group of the periodic
table). it is short by one election to the corresponding noble gas configuration, Helium (1s2). Hydrogen,
therefore, has resemblance to alkali metals, which lose one electron to form uni-positive ions, as well as with
halogens, whcih gain one electron to form uni-negative ion. Like alkali metals, hydrogen forms oxides,
healides and sulphides. However, unlike alkali metals, it has a very high ionization enthalpy and does not
possess metallic characteistics under normal conditions. In fact, in terms of ionization enthalpy hydrogen
resembles more with halogens, i H of Li is 520 kJ mol 1, H of F is 1680 kJ mol1 and i H of H

is 1312 kJ mol 1. Like halogens, if forms a diatomic molecule, combines with elements to form hydrides
and a large number of covalent compounds. However, in terms of reactivity, it is very low as compared to
halogens. It is always a matter of debate in which group hydrogen should be placed. It is best placed
separtely in the periodic table.

Isotpes of Hydrogen

Hydrogen has three isotopes : Protium (11 H), Deuterium (12 H) or D and Tritium (13 H) or T. These isotopes
differ from one another in respect of the presence of neutrons. Ordinary hydrogen, Protium, has no neu-
trons. Deuterium (also know as Heavy Hydrogen) has one and Tritium has two neutrons in the nucleus.
The predominant form is Protium. Terrestrial hydrogen contains 0.156% of Deuterium mostly in the form of
HD. The Tritium concentration is about one atom per 1018 atoms of Protium. Of these isotopes, only Tritium
is radioactiver and emits low energy Particles (t1/2 = 12.33 years).

Since the isotopes have the same electronic configuration, they have almost the same chemical properties.
The only difference is in their rates of reactions, mainly due to their different enthyalpy of bond dissociation.

Property Hydrogen Deuterium Tritium

Relative abundance (%) 99.985 0.0156 10-15
Relative atomic mass (g mol) 1.008 2.014 3.016
Melting point (K) 13.96 18.73 20.62
Boiling point (K) 20.39 23.67 25.0
Density (g L-1) 0.09 0.18 0.27

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Prepartion of Dihydrogen, H2
There are a number of methods for preparing dihydrogen from metals and metals and hydrides.
Laboratory Preparation of Dihydrogen
1. It is usually prepared by the reaction granulated zinc with dilute hydrochloric acid.

Zn  2H  
 Zn 2   H 2
2. It can also be prepared by the reaction of zinc with aqueous alkali.

Zn  2NaOH 
 Na 2 ZnO 2  H2
Sodium Zincate

Commercial Production of Dihydrogen :

1. Electroysis of acidified water using platinum electrodes gives hydrogen. This method gives very pure
H2 but it is very expensive

Electrolysis
2H2O 1 
 2H2  g   O2  g 
Traces of acid /base

2. It is obtained as a by product in the manufacture of sodium hydroxide and chlorine by the electrolysis
of brine solutionn. During electrolysis, the reactions that take place are :
At anode : 2Cl  aq  
 Cl2  g   2e

At cathode : 2H 2O  l   2e  
 H 2  g   2OH   aq 

The overall reaction is

2Na   aq   2Cl  aq   2H 2O  l  
 Cl2  g   H 2  g   2Na   aq   2OH   aq 

3. Reaction of steam on hydrocarbons or coke at high temperatures in the presence of catalyst yields
hydrogen.

1270K
 nCO   2n  1 H 2
Cn H 2n  2  nH 2O 
Ni
1270K
CH 4  g   H 2O  g  
 CO  g   3H 2  g 
Ni

The mixture of CO and H2 is called Water gas. As this mixture of CO and H2 is used for the synthesis
of methanol and a number of hydrocarbons, it is also called sythesis gas or ‘Syngas’.Nowadys‘syngas’
is produced from sewage, saw-dust, scrap wood, newspapers etc. The process of producing ‘synags’
from coal is called ‘Coal gasification’.
1270K
C  s   H 2O  g  
 CO  g   H 2  g 

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It is difficult to obtain pure H 2 from water gas, since CO is difficult to remove. Still CO may be liquified at
a low temperature under pressure. Thus seperating it from H 2 . One more thing that can be done is the gas
mixture can be mixed with steam, cooled to 400°C and passed over iron oxide to give H 2 and CO2 .

Fe
CO + H 2O  CO2 + H 2

This is called Water - gas shift reacrion. This reaction increases the amount of H2 and gives a method to
extract H2 easily. CO2 in mixture of CO2 and H2 can be removed by dissolving mixture in water under
pressure, or reacting mixture K2CO3 solution giving KHCO3, or by scrubbing mixture with sodium arsenite
solution.

Physical Properties :
 Dihydrogen is a colouless, odourless, tasteless, combustible gas. It is lighter than air and insoluble in
water.
Chemical Properties :
The chemical behaviour of dihydrogen (and for that mattter any molecule) is determined, to a large extent,
by bond issociation enthyalpy. The H-H bond dissociations enthalpy is the highest for a single bond between
two atoms any element. It is because of this factor that the dissociation of dihydrogen into its atoms is only
 0.081% around 2000K which increases to 95.5% at 5000K. Also, it is relatively inert at room tempera-
ture due to the high H-H bond enthalpy.
1. Reaction with halogens : It reacts with halogens, X2 to give hydrogen haliides, HX.
H 2  g   X 2  g  
 2HX  g   X  F, Cl, Br, I 
While the reaction with fluorine occurs even in the dark, with iodine it requires a catalyst.
2. Reaction with Dioxygen : It reacts with dioxygen to form water. The reaction is highly exothermic.
catalyst of heating
2H 2  g   O 2  g  
 2H 2O  l  ; H   285.9 kJ mol 1

3. Reaction with dinitrogen : It reacts with dinitrogen to form ammonia.

673K, 200atm
3H 2  g   N 2  g   2NH3  g  ; H   92.6 kJ mol1
Fe

This is the method for the manufacture of ammonia by the Haber’s process.
4. Reactions with metals : With many metals it combines at a high temperature to yield the corre-
sponding hydrides.

H 2  g   2M  g  
 2MH  s  ; where M is an alkali metal

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5. Reaction with metal ions and metal oxides : It reduces some metal ions in aqueous solution
and oxides of metals (less active than iron) into corresponding metals.

H 2  g   Pd 2   aq  
 Pd  s   2H  (aq)

yH 2  g   M x O y  s  
 xM  s   yH 2O 1

6. React ion w it h organic com pounds : It reacts with many organic compounds in the presence of
catalysts to give useful hydrogenated products of commercial importance. For example :
Hydrogenation of vegetable oils using nickel as catalyst give edible fats (margarine and vanaspathi
ghee).

Uses of Dihydrogen
 The largest single use of dihydrogen is in the synthesis of ammonia which is used in the manufacture of
nitric acid and nitrogenous fertilizers.
 Dihydrogen is used in the manufacture of vanaspati fat by the hydrogenation of polyunsaturated veg-
etable oils like soyabean, cotton seeds etc.
 It is used in the manufacture of bulk organic chemicals, particularly methanol.

cobalt
CO  g   2H 2  g   CH3OH  l 
catalyst
 It is widely used for the manufacture of metal hydrides.
 It is used a rocket fuel in space research.
 Dihydrgoen is used in fuel cells for generating electrical energy. If has many advantages over the
conventional fossil fuels and electric power. It does not produce any pollutions and releases greater
energy per unit mass of fuel in comparison to gasoline and other fuels.
Hydrides
Binary compounds of the elements with hydrogen are called hydrides. The type of hydride which an element
forms depends upon its electronegativity and hence on the type of bond formed. Hydrides are conveniently
studied under three classes.
(i) Ionic or salt like hydrides
(ii) Covalent or molecular hydrides
(iii) Metallic or interstitial hydrides
 Ionic or salt like hydrides : These are formed by metals of low electronegativity, i.e. alkali and
alkaline earth metals by direct reaction with H2 and some highly positive members of lanthanide series
with the exception of Be and Mg whose hydrides show significiant covalent character.

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Te stability of the hydrides decreases as the size of the cation increases.

LiH  NaH  KH  RbH  CsH

CaH 2  SrH 2  BaH 2

CaH2 is called Hydrolith.

 Covalent or Molecular hydrides : These hydrides are formed by all the true non-metals (except
zero group elements) and the elements like Al, Ga, Sn, Pb, Sb, Bi. Po, etc., which are normally
metallic in nature. The simple hydride of B and Ga are dimeric materials B2H6 (diborane) and Ga2H6
respectively and the hydride of aluminium is polymeric in nature, (AlH3)n.
 Metallic or interstitial hydrides : Many transition and inner-transition elements at elevated tem-
peratures absorb hydrogen into the interstices of their lattices to yield metal-like hydrides, often called
the interstitial hydrides. These hydrides are often non-stoichiometric and their composition vary with
temperature and pressure. Formulae of some of the hydrides of this class are :

TiH1.73 , CeH 2.7 , LaH 2.8 , PdH 0.60 , ZrH1.92

The interstitial hydrides have metallic appearance and their properties are closely related to those of
the parent metal. They posses reducing properties probably due to the presence of free hydrogen
atoms in the metal lattice.

Water
In the gas phase water is a bent molecule with a bond angle of 104.5o, and O-H bond length of 95.7 pm. It
is a highly polar molecule. Its orbital overlap is shown in figure below. In liquid phase, water molecules are
associated together by hydrogen bonds. In ice each oxygen atom is surrounded tetrahedrally by four other
oxygen atoms at a distance of 276 pm.

Chemical Properties of Water :

Water reacts with a large number of substances. Some of the important reactions are given below.
1. Amphoteric Nature : It has the ability to act as an acid as well as a base i.e., it behaves as an
amphoteric substance. In the Bronsted sense it acts as an acid with NH3 and a base with H2S.



H 2O  l   NH3  aq    
 OH  aq   NH 4  aq 


H 2O  l   H 2S  aq    
 H3O  aq   HS  aq 

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The auto-protolysis (self-ionization) of water takes places as follows :



H 2O  l   H 2O  l   
 H3O  aq   OH   aq 
acid 1 base  2 acid  2 base  1
 acid   base   conjugate acid   conjugate base

2. Redox Reactions Involving Water : Water can be easily reduced to dihydrogen by highly electroposi-
tive metals.

2H 2O  l   2Na  s  
 2NaOH  aq   H 2  g 

Water is oxidised to O2 during photosynthesis.

6CO 2  g   12H 2O  l  
 C6 H12O6  aq   6H 2O  l   6O 2  g 

With fluorine also it is oxidized to O2.

 4H   aq   4F  aq   O 2  g 
2F2  g   2H 2O  l  

3. Hydrolysis Reaction : Due to high dielctric constant, it has a very strong hydrating tendency. It dissolves
many ionic compounds. However, certain covalent and some ionic compound are hydrolysed in water.

P4O10  s   6H 2O  l  
 4H3PO 4  aq 

SiCl4  l   2H 2O  l  
 SiO2  s   4HCl  aq 

N3  s   3H 2O  l  
 NH 3  g   3OH   aq 

4. Hydrates Formation : From aqueous solutions many salts can be crystallised as hydrated salts. Water of
hydration are water molecules atttached to a compound that can be removed on heating. Such an associa-
tion of water is of different types viz.,
3
(i) Coordinated water e.g.,  Cr  H 2O   3Cl

(ii) Interstitial water e.g., BaCl2 . 2H 2O

2
(iii) Hydrogen-bonded wate e.e.,  Cu  H 2O 4  SO42  . H 2O in CuSO4 . 5 H 2O

Here in CuSO 4 . 5H 2O, four water molecules of hydration are coordinate bonded and one is hydro-
gen bonded.

Note : You will learn about hyrates later in Coordination Compounds.

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*Hard and Soft Water :

Rain water is almost pure (may contain some dissolved gases from the atmosphere). Being a good solvent,
when it flows on the surface of the earth, it dissolves many salts. Presence of calcium and magnesium salts
in the form of hydrogencarbonate, chloride and sulphate in water makes water ‘hard’. Hard water does not
give lather with soap. Water free from soluble salts of calcium and magnesium is called Soft water. It gives
lather with soap easily.
Hard water forms scum/precipitate with soap.Soap containg sodium strearate (C17H35COONa) reacts
with hard water precipitate out Ca/Mg strearate.

2 RCOONa(aq) + M 2+ (aq) 
 (RCOO)2 M  +2 Na + (aq) ; M is Ca/ Mg
Scum

e.g., 2C17 H35COONa  aq   Ca 2   aq  

  C17 H35COO 2 Ca   2Na   aq 
Scum
It is, therefore, unsuitable for laundry.
The hardness of water is of two types :
(i) Temporary hardness, and (ii) Permanent hardness.

Temporary Hardness :
Temporary hardness is due to the presence of magnesium and calcium hydrogen-carbonates. It can be
easily removed by :
1. Boiling : During boiling, the soluble Mg(HCO3)2 is converted into insoluble Mg(OH)2 and Ca(HCO3)2
changed to insoluble CaCO3. These precipitates can be removed by filteration.
Heating
Mg  HCO3 2 
 Mg  OH 2   2CO 2 
Heating
Ca  HCO3 2 
 CaCO3   H 2O  CO 2 

2. Clark’s method : In this method calculated amount of lime is added to hard water. It precipitates out
calcium carbonate and magnesium hydroxide which can be filtered off.

Ca  HCO3 2  Ca  OH  2 
 2CaCO3   2H 2O

Mg  HCO3 2  2Ca  OH 2 
 2CaCO3   Mg(OH)2   2H 2O

Note : Temporary hardness such as Ca(HCO3 )2 can also be removed by adding Na 2CO3 .

Ca(HCO3 )2 + Na 2CO3 
 CaCO3  +2 NaHCO3

Permanent Hardness :
It is due to the presence of soluble salts of magnesium and calcium in the form of chlorides and sulphates in
water. Permanent hardness is not removed by boiling. It can be removed by the following methods:

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1. Treatment with washing soda (Sodium carbonate) :

MCl2  Na 2CO3 
 MCO3   2NaCl (M  Mg, Ca)
MSO4  Na 2CO3 
 MCO3   Na 2SO4
(Washing soda)

2. Calogen’s method : Sodium hexametaphosphate (Na6P6O18), commercially called ‘Calgon’ or Graham’

salt’, when added to hard water, the following reactions take place :

 2Na   Na 4 P6O182 
Na 6 P6O18  M  Mg, Ca 
2
M 2   Na 4 P6O182  
  Na 2 MP6O18   2Na 

The complex anion keeps the Mg 2  and Ca 2  ions in solution.

3. Ion-exchange method : This method is also called zeolite/permutit process. Hydrated sodium aluminium
silicate is zeolite/permult prcess. For the sake of simplicity, sodium aluminum silicate (NaAlSiO4) can be
written as NaZ. When this is added in hard water, exchange reactins take place.

2NaZ(s)  M 2  (aq) 
 MZ2 (s)  2Na  (aq) (M  Mg, Ca)

Permutit/zeolite is said to be exhausted when all the sodium in it is used up. It is regenerated for further use
by treating with an aqueous chloride solution.

MZ2 (s)  2NaCl (aq) 

 2NaZ (s)  MCl2 (aq)
4. Synthetic Resins Methods : It is used in the prodcution of deionised water and more efficient than the
Zeolite process.
Water is passed through two differnt ion-exchange columns :
1st Column (Cation Exchange Coloumn) :
The resin exchange H  with Na  , Ca 2  and Mg 2 

resin  SO3H  M  
 resin  SO3M  H 
 
sulphonic acid
resin
2nd Coloumn (Anion Exchange Column) :
The resin exchanges OH  with Cl  , HCO3 , SO 42 

resin  NR 4 OH   X  
 resin NR 4 X   OH 
 
re sin with basic
group

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When all reactive sites on resins have been used they can be regenerated by treating first one with dil.
H 2SO 4 and second one with Na 2CO3 solution.

Heavy water, D2O

It is extensively used as a moderator in nuclear reactors and in exchange reactions for the study of reaction
mechanisms. It can be prepared by exhaustive electrolysis of water or as a by-product in some fertilizer
industries. It is used for the preparation of other deuterium compounds, for example :

 C2 D 2  Ca  OD 2
CaC2  2D 2O 

SO3  D2O 
 D 2SO 4
 3CD4  4Al  OD 3
Al4C3  12D 2O 

DEGREE OF HARDNESS Section - 6

Concentration of Solute in Terms of Parts per Million (or ppm) :

Concentration of solute ( in ppm) = mass of solute ( in gms) in 106 ml solution
It is used in determining the hardness of water which is due to the presence of bicarbonates (temporary
hardness), chlorides and sulphates (permanent hardness) of Calcium and Magnesium. Degree of Hardness
is defined as the number of parts of CaCO3 or equivalent to other calcium and magnesium salts present in

a million (106 ) parts of water..

Mass of CaCO3
Degree of Hardness   106 ppm
Mass of water

Illustrating the Concept :

How to calculate degree of hardness in a water sample containing 111 ppm of CaCl2 ?
100 111
E CaCO3 = = 50 ; E CaCl 2 = = 55.5
2 2
which means 50 gm of CaCO3  55.5 gm CaCl2
or 55.5 gm CaCl2  50 gm of CaCO3
 111.0 gm CaCl2  100 gm of CaCO3  100 ppm

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Illustration - 1Calculate the weight of CaO required to remove hardness of 106 L of water contain-
ing 1.62 gm of Ca(HCO3 )2 in 1.0 litre .

Solution :
Consider the reaction between CaO and Ca(HCO3 )2 .  moles of CaO required for 1.0 L of sample
= 0.01 (from stoichiometry)
CaO+ Ca(HCO3 )2 
 2 CaCO3 + H 2O
 moles of CaO required for 106 L of water
From stoichiometry, we have :
1 mole of Ca(HCO3 )2  1mole of CaO  0.01  106  104 moles.
 2 moles of CaCO3  grams of CaO  104  56  5.6  105 gm.
Now moles of bicarbonate in 1.0 L of sample
1.62
  0.01 [ M 0 of Ca(HCO 3 ) 2 = 162]
162

Illustration - 2 A particular water sample is found to contain 96.0 ppm of SO42  and 122.0 ppm of

HCO3 , with Ca 2  as the only cation. How many ppm of Ca 2  does this water contain ?
Solution :
CaSO 4  Ca 2+ + SO 24  96 ppm  96 g SO 24  in 106 mL H 2O

Ca(HCO3 ) 2  Ca 2+ + 2 HCO3  1.0 mol of SO 24   1.0 mol of Ca 2+

Every mol of CaSO 4 has equal mol of 122 ppm = 122 g HCO 3 in 10 6 mL H 2 O
Ca 2+ and SO 24  but every mole of
 2.0 mol of HCO3  1.0 mol of Ca 2+
Ca(HCO3 )2 has Ca 2+ half the mol of HCO3 .
Total Ca 2+ = 1+1 = 2.0 mol of Ca 2+

 80 g in10 6 mL H 2O
 Ca 2+  80 ppm

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Illustration - 3 A sample of hard water contains 1 mg CaCl3 and 1 mg MgCl3 per litre. Calculate the
hardness of water in terms of CaCO3 present in per 106 parts of water..

Solution :
50
55.5 gm CaCl2  50 gm CaCO3  1 mg CaCl2  mg CaCO3  0.9 mg CaCO3
55.5

47.5 gm MgCl2  50 gm CaCO3 50

 1mg MgCl2  mg CaCO3  1.05 mg CaCO 3
47.5

(0.9 +1.05)×103 gm
 Hardness in CaCO3 ppm   1.95ppm
1/1000

Illustration - 4 A sample of hard water contains 244 ppm of HCO3 ions. What is the minimum mass of
CaO required to remove ions completely from 1 kg of such water sample ?
Solution :
244 ppm HCO3  244 gm HCO3 in 1000 L  244 mg HCO3 in 1.0 L  4 mmoles HCO3 in 1.0 L
 2 mmoles Ca(HCO3 )2 in 1.0 L

 2CaCO3 + H 2O 
CaO + Ca(HCO3 )2   2 mmoles CaO in 1.0 L  2  56  112 mg CaO

Illustration - 5 250 ml of hard water is treated with 100 ml of 0.1 N Na2 CO3 to remove temporary
hardness. Excess of Na2 CO3 required 40 ml, 0.1 N HCl for complete neutralization. Calculate degree of
hardness of water.
Solution :
Meq of Na 2CO3 = meq of Ca(HCO3 )2 + meq of HCl
in hard water

Meq of CaCO3 formed = meq of Ca(HCO3 )2 = meq of Na 2CO3 + meq of HCl

in hard water

 100  0.1  40  0.1  6

6
mmoles of CaCO3 formed  3 (in factor of CaCO3 = 2 )
2

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Mass of CaCO3 formed  3  103  100  0.3 gm

Mass of hard water sample = 250 gm (Assuming density of water = 1 gm/ml)
Mass of CaCO3
Degree of Hardness  106 ppm
Mass of water
0.3
  106 ppm  1.2  103 ppm
250

IN-CHAPTER EXERCISE - E

1. How can the dihydrogen be obtained from coal gasification method ? How is its production enhanced ?
2. Arrange the following :
(i) CaH 2 , BeH 2 and TiH 2 in order of increasing electrical conductance.
(ii) LiH , NaH and CsH in order of increasing ionic character..
(iii) H - H, D - D and F – F in order of increasing bond dissolciation enthalpy.
(iv) NaH, MgH 2 and H 2O in order of increasing reducing property .

NOW ATTEMPT IN-CHAPTER EXERCISE-D FOR REMAINING QUESTIONS

NOW ATTEMPT OBJECTIVE WORKSHEET BEFORE PROCEEDING AHEAD IN THIS EBOOK

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SOLUTION TO IN-CHAPTER EXERCISE - D

1. NaHCO3  NaOH   Na 2CO3  H 2O

2. It is because hydration energy of group 1 hydroxides and carbonates is higher than lattice energy where as
in group 2 hydroxides and carbonates lattice energy dominates hydration energy.
3. K2CO3 can’t be prepared from soluay’s process because KHCO3 formed in the reaction is highly soluble
and hence can’t be separated from NH4Cl + KHCO3 mixture easily.
4. MgO and BeO are used for lining of steel making furnance because they are less reactive, have high melting
point, good conductivity towards heat and bad conductivity towards electricity.
5. Note : Read the question as this :
On the treatment with cold water, an element (A) reacted quietly, liberating a colourless, odourless gas (B)
and a compound (C). Gas (B) further reacts with element (A) to yield a solid product (D) which reacted
with water to give a basic solution (E). (E) is found to be same as (C). When carbon dioxide was bubbled
through solution (C) initially a white precipitate (F) is formed, but this redissolved forming solution (G) when
more CO2 was added. Precipitate (F) effervesced when moistened with conc. HCl acid and gave deep red
colouration to the burner flame. When (F) was heated at 1000°C, a white compound (H) was formed
which when heated with carbon at 1000°C gave a solid (I) of some commercial importance. Name the
substances (A) to (I).
SOLUTION :
(A) is calcium metal which reacts with water and evolves hydrogen (B) and Ca(OH)2 solution (C).
Ca  2H 2O 
 Ca(OH)2  H 2
“Colourless and Odourless gas”
(A) (C) (B)

Ca  H 2 
 CaH 2
Solid product
(B) (D)

CaH 2  2H 2O 
 Ca(OH) 2  2H 2
(E)
2  Ca(HCO ) CO
Ca(OH)2  CO2 
 CaCO3  3 2
(F) (G)

CaCO3  2HCl 
 CaCl2  H 2O  CO 2 “Gives deep red colouration to flame”
 C
CaCO3 
 CaO 
 CaC2
(F) (H) (I)

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SOLUTION TO IN-CHAPTER EXERCISE - E

1. Reaction of steam on coal at high temperatures in presence of catalyst yields hydrogen. This process is called
as coal gasification.
1270K
C(s)  H 2O(g)   CO(g)  H 2 (g)
Mixing this CO and H2 gas mixture with steam, cooling to 400°C and passing over iron oxide gives CO2 + H2
which finally increases the amount of H2.
Fe
CO  H2  CO2  H2
The above reaction is also called as water gas shift reaction.

2. (i) BeH2(covelent) < TiH2 < CaH2

(ii) LiH < NaH < CsH
(iii) F–F<H–H<D–D
(iv) H2O < MgH2 < NaH

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My Chapter Notes

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Illustration - 1

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