Ppt xii p block elements , By Dr Mona Srivastava

CLASS-12
SUBJECT- CHEMISTRY
CHAPTER - P BLOCK
ELEMENTS
Dr. Mona Srivastava
M.Sc. Ph. D. Chemistry

PERIOD 1
Learning objective –
GROUP 15 ELEMENTS –
 Introduction
 Occurrence
 Electronic configuration

P block elements in periodic table-

Introduction *Group 15 includes
* N, P, As, Sb , Bi
*2. As we go down
the group, there is a
shift from non-
metallic to metallic
through metalloid
character.
*N , P 2 non-metals,
*As , Sb 2metalloids
*Bi 1 metal.
*

Occurrence
Molecular nitrogen (N2) - It comprises of 78% by volume of
the atmosphere.
Ores of nitrogen-
 1. Sodium nitrate, NaNO3 (called Chile saltpetre and )
2.Potassium nitrate (Indian saltpetre).
 3. Proteins - It is found in the form of proteins in plants and
animals.
Phosphorus - occurs in minerals.
1.Fluorapatite - Ca9 (PO4)6. CaF2
(the main components of phosphate rocks.)
 Bones/ Living cells – P is an essential constituent of animal and
plant matter. It is present in bones as well as in living cells.
Milk/eggs -Phosphoproteins are present in milk and eggs.
Arsenic, antimony and bismuth are found mainly as
sulphide minerals.

Their valence shell electronic
configuration is ns2np3
Electronic configuration

*Atomic and ionic radii
*Covalent and ionic (in a particular state) radii
increase in size down the group.
*There is a considerable increase in covalent radius
from N to P.
*However, from As to Bi only a small increase in
covalent radius is observed.
*This is due to the presence of completely filled
d and/ f orbitals in heavier members.

*Ionization Enthalpies
*Ionization enthalpy decreases down the group due
to gradual increase in atomic size.
*The ionization enthalpy of the group 15 elements is
much greater than that of group 14 elements – ( In
period)
*Reason- Because of the extra stable half-filled p
orbitals electronic configuration and smaller size,
*The order of successive ionization enthalpies, as
expected is
 H1 <  H2 <  H3 .

Electronegativities
The electronegativity value, in general,
decreases down the group with
increasing atomic size. But amongst
the heavier elements, the difference is
not that much pronounced

*Period 2
Learnin objectives -
Physical properties
Oxuidation states
Anomelous properties of Nitrogen

Physical Properties
1. All the elements of this group are polyatomic.
2. Dinitrogen is a diatomic gas while all others are solids.
3. Metallic character increases down the group.
4. Nitrogen and phosphorus are non-metals.
5. Arsenic and antimony metalloids .
6. Bismuth is a metal.
7. This is due to decrease in ionisation enthalpy and
increase in atomic size.
8. The boiling points, in general, increase from top to
bottom in the group .
9. The melting point increases upto arsenic and then
decreases upto bismuth.
10. Except nitrogen, all the elements show allotropy.

Oxidation states and trends in chemical
reactivity
All the elements of group 15 have 5 electrons in their
outermost orbit. They need only 3 electrons to complete
their octet configuration. The octet can be achieved either by
gaining 3 electrons or by sharing 3 electrons by means of
covalent bonds. As a result, the common negative oxidation
state of these elements is - 3.
As we move down the group, the tendency to exhibit -3
oxidation state decreases. This is due to the increase in
atomic size and metallic character.
Group 15 elements also show positive oxidation states of +3
and +5 by forming covalent bonds. Due to the inert pair
affect the stability of +5 oxidation state decreases down the
group, while that of +3 oxidation state increases.
contd…….

Continued ……
Nitrogen has only s- and p-orbitals, but no d-orbitals
in its valance shell. Therefore, nitrogen can show a
maximum covalency of 4.
 A covalency of four is obtained by sharing its lone
pair of electron with another atom or ion.
Phosphorus and the remaining elements can
exhibit a covalency of five and a maximum covalency,
also called expanded covalency of six.
This is possible because of the presence of vacant d-
orbitals in the valence shell.
All the compounds of group fifteen elements, which
exhibit a +5 oxidation state, are covalent.

*Nitrogen differs from the rest of the
members of this group due to
1. Small size 2. non-availability of d orbitals.
3. high ionisation enthalpy 4. high electronegativity
*Nitrogen has unique ability to form -
* pp-pp multiple bonds with itself and with other
elements (C, O). Thus, nitrogen exists as a diatomic
molecule with a triple bond (one s and two p)
between the two atoms.
*Heavier elements of this group do not form pp-pp
bonds because of their large atomic orbitals due to
which they cannot have effective overlapping.
Contd……

*
* Phosphorus, Arsenic And Antimony -
Form single bonds as P–P, As–As and Sb–Sb
smuth forms metallic bonds in elemental
state.
N–N bond is weaker than the single P–P
bond .
Reason- The N-N single bonds have high
interelectronic repulsion of the non-
bonding electrons, owing to the small bond
length. As a result the catenation tendency
is weaker in N2 , forming weaker covalent
bonds.

*Absence of d orbitals in valence shell in N2 --
* Due to absence of d orbitals N2 has unusual
chemistry than rest of the members.
*Nitrogen cannot form dp –pp bond as the heavier
*elements can e.g., R3P = O or R3P = CH2 (R = alkyl
group).
*P and As can form dp –dp bond also with
transition metals if their compounds like P(C2H5)3
and As(C6H5)3 act as ligands.

Period 3
*Learning objectives-
Chemical reactivity of nitrogen towards-
Hydrogen
Oxygen
Halogen

Reactivity towards hydrogen
*All the elements of Group 15 form hydrides of the type EH3
* E = N, P, As, Sb or Bi.
* The hydrides show regular variation in their
properties.
*The stability of hydrides decreases from NH3 to BiH3
*The reducing character of the hydrides increases.
*Ammonia is only a mild reducing agent while BiH3 is
the strongest reducing agent of all the hydrides.
*Basicity also decreases in the order
NH3 > PH3 > AsH3 > SbH3 > BiH3.

*Reactivity towards oxygen:
*All elements of 15 group form two types of oxides:
* E2O3 and E2O5.
*The oxide in the higher oxidation state of the
element is more acidic than that of lower oxidation
state.
* Their acidic character decreases down the group.
* E2O3 - Nitrogen and phosphorus - purely acidic,
*E2O3 - Arsenic and antimony - amphoteric
*E2o3 - Bismuth - Basic.

These elements of 15 group form form types
of halides: 1. EX3 , 2. EX5
Nitrogen does not form pentahalide due to
non-availability of the d orbitals in its
valence shell.
Pentahalides are more covalent than
trihalides.
 All the trihalides of these elements except
those of nitrogen are stable.
In N2 only NF3 is known to be stable.
Trihalides are covalent in nature except
BiF3.

All the elements of group 15 react with metals
to form binary compounds exhibiting –3
oxidation state.
Ex. Ca3N2 (calcium nitride) , Ca3P2 (calcium
phosphide),
Na3As2 (sodium arsenide), Zn3Sb2 (zinc
antimonide)
 Mg3Bi2 (magnesium bismuthide).

Q.1. Though nitrogen exhibits +5 oxidation state, it
does not form pentahalide. (Give reason)
Ans. - Nitrogen with n = 2, has s and p orbitals
only. It does not have d orbitals to expand its
covalence beyond four. That is why it does not form
pentahalide.
 Q.2.Give reason - PH3 has lower boiling point than
NH3. Why?
 Ans.-Unlike NH3, PH3 molecules are not
associated through hydrogen bonding in liquid
state. That is why the boiling point of PH3 is lower
than NH3.

Assignment No. 3
Q.1. Why are pentahalides more
covalent than trihalides ?
Q.2. Why is BiH3 the strongest reducing
agent amongst all Group 15 elements ?

*Period 4
* Learning objectuve –
Dinitrogen –
A. commercial production
Physical properties
Chemical pproperties
Uses.

DiNitrogen ( N2)
•Introduction-
•Nitrogen was discovered in 1770 by Scheele and
Priestley.
•This nonmetallic element has no color, taste or odor
and is present in nature as a noncombustible gas.
•When compared with the rest of Group 15, nitrogen
has the highest electronegativity which makes it the
most nonmetallic of the group.
•The common oxidation states of N2 are +5, +3, and -3.
• N2 makes up about 0.002% of the earth's crust.
• It constitutes 78% of the earth’s atmosphere by
volume.

Commercial production-
N2 is produced commercially by –
“The liquefaction and fractional distillation of air.
Liquid dinitrogen (b.p. 77.2 K) distils out first leaving
behind liquid oxygen (b.p. 90 K).”

Laboratory method
Dinitrogen in laboratory is prepared by treating -
1. Ammonium chloride (aq) + sodium nitrite.
NH4CI(aq) + NaNO2(aq) N2(g) + 2H2O(l) + NaCl (aq)
Small amounts of NO and HNO3 are also formed in this
reaction; these impurities can be removed by passing the
gas through aqueous sulphuric acid containing potassium
dichromate.
2. It can also be obtained by the thermal decomposition of
ammonium dichromate.
Heat
(NH4)2Cr2O7 N2 + 4H2O + Cr2O3

*
1. N2 is a colorless, odorless, tasteless and non-toxic gas.
2. Nitrogen atom has two stable isotopes: 14N and 15N.
3. It has a very low solubility in water (23.2 cm3 per litre of water at 273 K
and 1 bar P )
4.Low freezing and boiling points .

N2 is inert at room temperature because of the high bond
enthalpy of N N bond.
Reactivity, increases rapidly with rise in temperature.
 1. Reaction with metals and nonmetals ( at higher temp)
a. N2 reacts some metals to form ionic nitrides
6Li + N2 2Li3N 3Mg + N2
b. Non-metals to form covalent nitrides.
3 Mg + N2 Mg3N2
 2. Reaction with hydrogen-
t combines with hydrogen at about 773 K in the presence of
catalyst (Haber’s Process) to form ammonia:
N2(g) + 3H2(g) 2NH3(g) f H 0
= – 46.1 kJmol–1

Chemical Properties - Continued
*

Period -5
Learning objectives-
Ammonia
 Manufacture
 Physical properties
 Chemical properties
 Uses

NH3
*Preparation –
*1. Formation in air- NH3 is present in small quantities in
air and soil , it is formed by the decay of nitrogenous organic
matter
* e.g., urea. NH2CONH2 + 2H2O NH4 2CO3
* NH4 2CO3 ⇌ 2NH3 + H2O + CO2
*2. On a small scale – NH3 in small scale obtained from
ammonium salts by their decomposition followed by further
treatment with caustic soda or calcium hydroxide.
*2NH4Cl + Ca(OH)2 2NH3 + 2H2O + CaCl2
(NH4)2 SO4 + 2NaOH 2NH3 + 2H2O + Na2SO4

* Commercial preparation of ammonia
*Commercially NH3 is manufactured by Haber’s process.
* N2(g) + 3H2(g) ⇌ 2NH3(g) f H0 = – 46.1 kJ mol–1

Conditions favoring formation of NH3
As per le Chatelier’s principle, following conditions
woulf favour the formation of ammonia-
1. High pressure would favour the formation of
ammonia.
 ( 200 × 105 Pa (about 200 atm),.
2. A temperature of 400-450 0 C ( 700 K) helps in
reaction to go in forward direction.
3. The use of a catalyst such as iron oxide.
4. Catalytic promotor - Small amounts of K2O and
Al2O3 to increase the rate of attainment of equilibrium.
 (The flow chart for the production of ammonia is shown
in previous fig. , in which its shown that earlier iron
was used as a catalyst with molybdenum as a promoter.)

Physical properties of ammonia
* 1.Ammonia is a colourless gas with a pungent odour.
*2. Its freezing and boiling points are 198.4 and 239.7 K .
*3. In the solid and liquid states, it is associated through hydrogen
bonds and therefore it accounts for its higher melting and boiling
points.
*Geometry -
* 1. The geometry of the molecule is
* trigonal pyramidal with the
*nitrogen atom at the apex.
*2. It has three bond pairs and
*one lone pair of electrons as
* shown in the structure.

Chemical Properties of ammonia
*1. Highly soluble in water- NH3 gas is highly soluble in water. Its
aqueous solution is weakly basic due to the formation of OH– ions.
* NH3(g) + H2O(l) ⇌ NH4+
(aq) + OH– (aq)
*2. Formation of ammonium salts- It forms ammonium salts with
acids, e.g., NH4Cl, (NH4)2 SO4, etc.
*As a weak base, these salts sprecipitate the hydroxides
(hydrated oxides in case of some metals) of many metals from their
salt solutions. For example,
*ZnSO4 aq + 2NH4OHaq Zn (OH)2 (s) + (NH4) 2 SO4 aq
(white )ppt

Chemical properties continued…….
*The presence of a lone pair of electrons on the nitrogen atom of the
ammonia molecule makes it a Lewis base.
*It donates the electron pair and forms linkage with metal ions
and the formation of such complex compounds finds applications in
detection of metal ions such as Cu2+, Ag+:
* Ex. Cu2+ (aq) + 4 NH3(aq) ⇌ [Cu(NH3)4]2+ (aq)
* (blue) (deep blue)
* Ag +(aq) + Cl- (aq) AgCls (s)
(colourless) (white ppt)
*AgCls + 2NH3 (aq) [Ag (NH3)2] Cl (aq)
* (white ppt) (colourless)

Uses of Ammonia
*1. Ammonia is used to produce various nitrogenous
fertilisers -
Ex. 1. ammonium nitrate
2. urea,
3. ammonium phosphate
4. ammonium sulphate)
*2. In the manufacture of some inorganic nitrogen
compounds, the most important one being nitric acid.
*3. Liquid ammonia is also used as a refrigerant.

Assignment No. 5
*Q.1. Why does nitrogen show catenation properties
less than phosphorous ?
*Q.2. How is ammonia manufactured industrially?
*Q.3.Discuss the trends in the chemical reactivity of
group 15 elements.
*Q.3.Mention the conditions required to maximize the
yield of ammonia.
*Q.4. How does ammonia react with a solution of
Cu2+?

Period 6
Learning objective
Nitric acid
 Preparation
 Properties
 uses

Nitric acid (HNO3)
Nitrogen forms 3 types of oxoacids
1. H2N2O2 (hyponitrous acid).
2.HNO2 (nitrous acid)
3.HNO3 (nitric acid).
4. Amongst them HNO3 is most important.

Preparation-
• Laboratory method - In the laboratory, nitric acid is prepared
• NaNO3 + H2SO4 (conc) Heating NaHSO4 + HNO3
Or KNO3 in a in a glass retort
• 2. Large scale(commercial) production –By Ostwald’s
process.
• Method- By the catalytic oxidation of NH3 using atmospheric O2.
• 4NH3 (g)+ 5O2 (g)  4NO(g)+ 6H2O(g) (from air)
• Nitric oxide thus formed combines with oxygen giving NO2.
• 2NOg + O2 g ⇌ 2NO2 g
• Nitrogen dioxide so formed, dissolves in water to give HNO3.
• 3NO2(g) + H2O (l)  2HNO3 (aq)+ NO(g)
• NO thus formed is recycled and the aqueous HNO3 can be
concentrated by distillation upto ~ 68% by mass.
• Concentration upto 98% can be achieved by dehydration with
concentrated H2SO4.

Ostwald's method of preparation of HNO3
Catalytic oxidation of ammonia by air-

Properties of Nitric acid
Physical properties-
1. It is a colorless liquid (f.p. 231.4 K and b.p. 355.6 K).
 Laboratory grade HNO3 has ~ 68% of the HNO3 by mass.
 It has a specific gravity of 1.504.
Structure- In the gaseous state, HNO3 exists as a
planar molecule with the structure as shown-
Chemical properties – 1. In aqueous solution, nitric acid
behaves as a strong acid giving hydronium and nitrate ions.
HNO3 (aq) + H2O (l)  H3O+(aq) + NO3
-
(aq)
2. Concentrated nitric acid is a strong oxidizing agent .
It attacks most metals except noble metals such as gold and
platinum.

Chemical Properties Continued-
3. The products of oxidation depend upon the -
1. Concentration of acid 2. Temperature
3. The nature of the material undergoing oxidation.
3Cu + 8 HNO3 (dilute)  3Cu(NO3)2 + 2NO + 4H2O
Cu + 4HNO3 (conc.)  Cu(NO3)2 + 2NO2 + 2H2O
4. Reaction of nitric acid on Zinc-
a. With dil. dilute nitric acid -
4Zn + 10HNO3(dilute)  4 Zn (NO3)2 + 5H2O + N2O
b. With conc. nitric acid -
Zn + 4HNO3(conc.)  Zn (NO3)2 + 2H2O + 2NO2
Some metals (e.g., Cr, Al) do not dissolve in concentrated
nitric acid because of the formation of a passive film of
oxide on the surface.

5. Reaction concentrated nitric acid with non–metals and
their comp. – ( oxidation takes place)
I2 + 10HNO3 2HIO3 + 10NO2 + 4H2O
C + 4HNO3 CO2 + 2H2O + 4NO2
S8 + 48HNO3  8H2SO4 + 48NO2 + 16H2O
P4 + 20HNO34H3PO4 + 20NO2 + 4H2O
6.Brown Ring Test: Its performed for the detection of nitrates in solu.
The test is usually carried out in two steps-
Step.1-Adding dilute FeSO4 solu. to an aq. solu. containing nitrate ion.
Step no.2. - Then carefully adding concentrated sulphuric acid along
the sides of the test tube.
A brown ring at the interface between the solution and H2SO4
acid layers indicates the presence of nitrate ion in solution.
NO3-
+ 3Fe2+ + 4H+ NO + 3Fe3+ + 2H2O
[Fe(H2O)6]2 + NO [Fe (H2O)5 (NO)]2+ + H2O (brown)

Uses:
 1. In the manufacture of ammonium nitrate for
fertilizers and nitrates.
 2. In manufacture of explosives and pyrotechnics.
 3. Used for the preparation of nitroglycerin,
trinitrotoluene and other organic nitro compounds.
4. In the pickling of stainless steel, etching of metals
and as an oxidizer in rocket fuels.

Period – 7
Learning objective
Group 16 elements –
Abundance
Electronic configuration
Atomic size and ionic radii
Ionization enthalpy
Electron gain enthalpy
Electronegativity

Group 16 elements (Chalcogens)
Oxygen ( O)
Sulphur (S)
Selenium(Se)
Tellurium ( Te)
Polonium (Po)
Livermortium ( Lv)

Occurrence -
Occurrence –
1. Oxygen is the most abundant of all the elements on earth.
2. Oxygen forms about 46.6% by mass of earth’s crust. Dry air
contains 20.946% oxygen by volume.
3. Sulphur in the earth’s crust is only 0.03-0.1%
4. Compounds of S -
1. As Sulphates- a. CaSO4.2H2O (gypsum )
b. MgSO4.7H2O ( Epsom salt)
c. BaSO4 (baryte)
2. As Sulphides – 1. PbS (galena)
2. ZnS zinc (blende )
3. CuFeS2 (copper pyrites)
5. Se / Te are also found as metal selenides , tellurides in sulphide
ores.
6. Po occurs in nature as a decay product of thorium and
uranium minerals.

Periodic properties
1. Electronic configuration- The elements of Group16 have 6
electrons in the outermost shell and have ns2np4 general
electronic configuration.
2. Atomic and Ionic Radii - Due to increase in the number of
shells, atomic and ionic radii increase from top to bottom in the
group. The size of oxygen atom is, however, exceptionally
small.
3. Ionizations Enthalpy –
a. Ionization enthalpy decreases down the group due to
increase in size.
b. However, the elements of this group have lower ionization enthalpy
values compared to those of Group15 in the corresponding periods.
Reason - Group 15 elements have extra stable half-filled p
orbitals electronic configurations.
Continued ……….

Continued
4. Electron gain enthalpy –
a. Oxygen has a compact
atom, therefore it has less
negative electron gain enthalpy
than sulphur.
b. But for the elements
from sulphur onwards the
value again becomes less
negative upto polonium.

Electronegativity
*After fluorine, oxygen has the highest
electronegativity value amongst the
elements.
*Within the group, electronegativity
decreases with an increase in atomic
number.
* Due to this the metallic character increases
from oxygen to polonium.

Figure - Trends of periodic properties-

Physical properties
1. The group 15 has –
a. Non metals - Oxygen and sulphur , polonium
b. Metalloids - Selenium and tellurium
c. Metal – Polonium ( radioactive / short lived metal
with a half life 13.8 days).
2. All these elements exhibit allotropy.
3. The melting and boiling points increase with an
increase in atomic number down the group.
4. The large difference between the melting and boiling
points of oxygen and sulphur may be explained on the
basis of their atomicity;
a. Oxygen exists as diatomic molecule (O2)
b. Sulphur exists as polyatomic molecule (S8).

1. Oxidation states and trends in chemical reactivity –
a. The stability of -2 ox. state decreases down the group.
b. Po - hardly shows –2 oxidation state.
c. O2- Shows a negative ox. State of -2 due to its
electronegativity. ( except in OF2 with ox. state is + 2 )
2. S , Se and Te – of the group exhibit + 2, + 4, + 6 oxidation
states but + 4 and + 6 are more common.
Ex- S , Se and Te –
+ 4 oxidation state - in their compounds made with O2 .
+ 6 oxidation state - with their compounds made with F2.
3. The stability of + 6 oxidation state decreases down the
group and stability of + 4 oxidation state increase (inert
pair effect).
4. Bonding in +4 and +6 oxidation states are primarily
covalent.

*Reasons for anomalous behavior of O2 -
1. It has a very small size and high electronegativity.
2. Absence of d orbitals (non availability of valence
electrons in d orbital)
Anomalous behavior shown by O2 as follows -
1. Oxygen is diatomic whereas others are polyatomic.
2. Oxygen is a gas while others are solid.
3. Oxygen shows negative oxidation states except in OF2
4. Oxygen is paramagnetic while others are diamagnetic.
5. Oxygen shows presence of strong hydrogen bonding in H2O
which is not found in H2S.
6. Due to the absence of d orbitals in oxygen limits its
covalency to 4 and rarely exceeds two. On the other hand, in
case of other elements of the group, the valence shells can be
expanded and covalence exceeds four.

1. All the elements of Group 16 form hydrides of the type
H2E . Where (E = S, Se, Te, Po).
2. Their acidic character increases from H2O to H2Te.
Reason- The increase in acidic character is due to the
decrease in bond (H–E) dissociation enthalpy down the
group. The thermal stability of hydrides depends on
bond (H–E) dissociation enthalpies. Therefore the
thermal stability of hydrides also decreases from H2O
to H2Po.
3. All the hydrides except water possess reducing
property and this character increases from H2S to H2Te

1. Elements of Group 16 form a large number of halides –
2. Ex. - EX6 , EX 4 and EX 2 where E is an element , X is a halogen.
3. The stability of the halides –Decreases down the gr. F–>Cl–>Br –> I–
4. Tetrafluorides- 1.SF4 is a gas, 2. SeF4 a liquid , 3.TeF4 a solid.
5. Fluorides - 1.Have sp3d hybridisation have trigonal
bipyramidal structures with lone pair of electrons. ( see-saw
geometry.)
6. Formation of dichlorides – Except Se all elements form
dichlorides and dibromides which have sp3 hybridisation
(tetrahedral structure) .
7. The well known monohalides are dimeric in nature. Examples are
S2F2, S2Cl2, S2Br2, Se2Cl2 and Se2Br2. These dimeric halides undergo
disproportionation as given below:
2Se2Cl2 → SeCl4 + 3Se

Period 8
Dioxygen –
 Preparation
 Manufacure
 Properties
 Uses

Laboratory method-
1. By heating oxygen containing salts such as chlorates, nitrates and
permanganates.
 2KClO 3 Heat 2KCl + 3O2
* MnO2
2. By the thermal decomposition of the oxides-
Ex 1.metals low in the electrochemical series ,
2. higher oxides of some metals.
 2Ag2O(s) → 4Ag(s) + O2(g)
 2Pb3O4(s) → 6PbO(s) + O2(g)
 2HgO(s) → 2Hg(l) + O2(g)
 2PbO2(s) → 2PbO(s) + O2(g)
3. By the rapid decomposition of Hydrogen peroxide - By catalysts
such as finely divided metals and manganese dioxide.
 2H2O2(aq) → 2H2O(1) + O2(g)

Commercial production of O2 -
By the electrolysis of water - On large scale it can be
prepared from water by its electrolysis .
At anode – oxygen is released.
At Cathod – Hydrogen is released .
Industrial Method - ( from air)
1. First removal of carbon dioxide and water vapour
from air is done.
2. Then, the remaining gases are liquefied and
fractionally distilled to give dinitrogen and dioxygen.

Physical Properties-
1. O2 colorless and odourless gas.
2. Its solubility in water is 3.08 cm3 in 100 cm3
water at 293 K which is sufficient for the support
of marine and aquatic life.
3. It liquefies at 90 K and freezes at 55 K.
4. It has three stable isotopes: 16O, 17O and 18O.
5. Molecular oxygen, O2 is unique in being .
paramagnetic inspite of having even number of
electrons.

Chemical properties of O2 are follows –
1. Dioxygen directly reacts with nearly all
metals and non-metals except some noble
metals and gases ( Au, Pt).
2. Its combination reaction with other
elements is often strongly exothermic for
which external heating is required as bond
dissociation enthalpy .
(oxygen - oxygen double bond is 493.4 kJ
mol–1).

Some important reactions of dioxygen with –
a. with metals –
2Ca + O 2 → 2CaO
 4Al + 3O 2 → 2Al 2 O3
2ZnS + 3O2 → 2ZnO + 2SO2
b. non-metals
P4 + 5O2 → P4O10
 C+O2 →CO2
c. Some compounds
 CH4 + 2O2 → CO2 + 2H2O
 2SO2 + O2
V2O5 → 2SO3
4HCl + O 2
CuCl 2 2Cl 2 + 2H 2 O

1. Most important use of oxygen is in respiration and
combustion .
2. processes, oxygen is used in oxyacetylene welding,
3. In the manufacture of many metals, particularly steel.
4. Oxygen cylinders are widely used in hospitals, high
altitude flying and in mountaineering.
5. The combustion of fuels, e.g., hydrazines in liquid
oxygen, provides tremendous thrust in rockets.

Assignment- 8
Q.1.Which of the following does not react with oxygen
directly?
Zn, Ti, Pt, Fe
Q.2.Complete the following reactions:
C2H4 + O2 →
4Al + 3 O2 →

Learning objective-
 Oxides
Types of oxides
Ozone
Preparation
Properties
uses

 Some metallic oxides exhibit a dual behaviour.
They show characteristics of both acidic as well as basic oxides. Such
oxides are known as amphoteric oxides.
They react with acids as well as alkalies.
known as neutral oxides.
• Examples of neutral oxides are CO, NO and N2O.
For example, Al2O3 reacts with acids as well as alkalies.
• Al2O3( s ) + 6HCl(aq) + 9H 2O (l ) → 2 [ Al(H2O)6 ]3+( aq ) + 6Cl− ( aq)
• Al 2O3(s) + 6NaOH ( aq ) + 3H2O(l) → 2Na3[ Al ( OH)6 ](aq ).

Occurrence –
 Ozone is an allotropic form of oxygen.
 It is not found at sea level too being reactive .
 At a height of about 20 kilometers, it is formed from
atmospheric oxygen in the presence of sunlight.
 This ozone layer protects the earth’s surface from an
excessive concentration of ultraviolet (UV) radiations.
Preparation –
When a slow dry stream of oxygen is passed through a
silent electrical discharge, conversion of oxygen to
ozone (10%) occurs. The product is known as ionized
oxygen –
3O2 → 2O3 HV (298 K) = +142 kJ mol–1

Structure-
The two oxygen-oxygen bond lengths in the ozone
molecule are identical (128 pm) and the molecule is
angular as expected with a bond angle of about 117o. It
is a resonance
hybrid of two main forms :

Physical –
1. Pure ozone is a pale blue gas, dark blue liquid and violet-
black solid.
2. Ozone has a characteristic smell and in small
concentrations it is harmless.
3. However, if the concentration rises ≥ 100 parts per
million, it creats breathing problem , headache and
nausea.
Chemical Properties –
1. Thermal decomposition –
Ozone is thermodynamically unstable
2O3 → 3O2 (exothermic, -G )
Due to large negative G values , it gets converted into O2 , and
the large concentration of O3 in one place which can be
explosive can be prevented .

2. Powerful oxidising agent- It acts as powerful oxidising agent due
to the ease with which it liberates atoms of nascent oxygen
(O3 → O2 + O).
1. It oxidises lead sulphide to lead sulphate and iodide ions to iodine.
1. PbS(s) + 4O3(g) → PbSO4(s) + 4O2(g)
2. 2I–(aq) + H2O(l) + O3(g) → 2OH–(aq) + I2(s) + O2(g)
3. Quantitative estimation – When O3 reacts with an excess of KI
at pH 9.2 , I2 is liberated which can be titrated against a standard
solution of sodium thiosulphate and its quantity can be found out.
Threats to O3 Layer –
a. From exhaust of supersonic jets carrying NO - There is a very fast
reaction between - NO (g ) + O3 (g ) → NO2 (g ) + O 2 (g)
The exhaust of jets carrying may deplete the layer of O3.
b. Another threat to this ozone layer is probably posed by the use of
freons (CFC)which are used in aerosol sprays and as refrigerants.
*

Threats to O3 Layer –
a. From exhaust of supersonic jets carrying NO - There
is a very fast reaction between -
NO (g ) + O3 (g ) → NO2 (g ) + O 2 (g)
The exhaust of jets carrying NO may deplete layer of O3.
b. Another threat to this ozone layer is probably posed
by the use of freons (CFC)which are used in aerosol
sprays and as refrigerants.
Uses – 1. It is used as a germicide, disinfectant and for
sterilising water.
2. It is also used for bleaching oils, ivory, flour, starch, etc.
3. 3. It acts as an oxidising agent in the manufacture of
potassium permanganate.

Assignment- 9
Q.1. Why does O3 act as a powerful oxidising agent?
Q.2. How is O3 estimated quantitatively ?
Q.3. What are the reasons of ozone layer depletion?

Learning Objective-
Sulphur (allotropic forms)
Sulphur Dioxide
Preparation
Properties
Uses

Sulphur (allotropic forms)
Exists mainly in 2 alloptropic forms-
(Β)
(α )

Facts-
1. Both rhombic and monoclinic sulphur have S8 molecules.
2. These S8 molecules are packed to give different crystal
structures.
3. The S8 ring in both the forms has a crown shape.
4. The molecular dimensions are given in pictures.
5. α Sulphur  β Sulphur ( at 369 K)

Q. Which form of sulphur shows paramagnetic behaviour
Ans - In vapour state sulphur partly exists as S2 molecule which has 2
unpaired electrons in the antibonding π * orbitals like O2 and, hence,
exhibits paramagnetism.
*Sulphur Dioxide – Preparation
By burning Sulphur -
S(s) + O2(g) → SO2 (g) [ note- 6-8 % SO3 is also formed.]
Laboratory method - In the laboratory it is readily generated by
treating a SO3-- with dil. H2SO4
SO3
2-(aq) + 2H+ (aq) → H2O(l) + SO2 (g)
Industrial Method -
Industrially, it is produced as a by-product of the roasting of
sulphide ores.
4FeS2 ( s ) + 11O2 (g ) → 2Fe2 O3 ( s) + 8SO2 (g )
* The gas after drying is liquefied under pressure and stored in
steel cylinders

Sulphur dioxide – [ SO2]
Physical Properties -
1. It is a colourless gas with pungent smell highly soluble in water.
2. It liquefies at room temp. under a P=2atm and boils at 263 K.
Chemical Properties -
1. Forms sulphurous acid when dissolved in H20.
SO2 ( g) + H2 O (l) H2 SO3 (aq)
2. It reacts readily with sodium hydroxide solution, forming sodium
sulphite, which then reacts with more sulphur dioxide to form sodium
hydrogen sulphite.
*2NaOH + SO2 → Na2SO3 + H2O
* Na2SO3 + H2O + SO2 → 2NaHSO3
3. In its reaction with water and alkalies, the behaviour of sulphur
dioxide is very similar to that of carbon dioxide.
Contd…………..

Chemical properties contd…..
4. Reaction with Cl2 - SO2 reacts with chlorine in the presence of
charcoal (which acts as a catalyst) to give sulphuryl chloride, SO2Cl2. It
is oxidised to sulphur trioxide by oxygen in the presence of
vanadium(V) oxide catalyst.
SO2(g) + Cl2 (g) → SO2Cl2(l)
V2O5
2SO2 (g) + O2 (g) 2SO3 (g)
5. Moist, sulphur dioxide behaves as a reducing agent.
For example, it converts iron(III) ions to iron(II) ions and
decolourises acidified potassium permanganate(VII) solution; the
latter reaction is a convenient test for the gas.
2Fe3 + + SO2 + 2H2 O → 2Fe2 + + SO2
4
− + 4H+
5SO2 + 2MnO−
4 + 2H2 O → 5SO2
4
− + 4H+ + 2Mn2+

Uses of SO2
Sulphur dioxide is used -
 In refining petroleum and sugar
 In bleaching wool and silk and
 As an anti-chlor, disinfectant and preservative.
 Sulphuric acid, sodium hydrogen sulphite and
calcium hydrogen sulphite (industrial chemicals)
are manufactured from sulphur dioxide.
 Liquid so2 is used as a solvent to dissolve a number
of organic and inorganic chemicals.

Structures of some important oxoacids of
sulphur

Assignment no. 10
Q.1 . What happens when sulphur dioxide is passed
through an aqueous solution of Fe(III) salt?
Q.2. Comment on the nature of two S–O bonds formed
in SO2 molecule. Are the two S–O bonds in this
molecule equal ?
Q.3. How is the presence of SO2 detected ?

Learning objective
Group 17 elements-
Occurrence
Electronic configuration
Atomic and ionic radii
Ionization enthalpy
Electron gain enthalpy
Electronegativity

Occurrence
1. F2 and Cl2 fairly abundant while Br2 and I2 less abundant .
2. F2 is present mainly as insoluble fluorides .
Ex. 1. Fluorspar [ CaF2 ] 2. Cryolite [ Na3AlF6 ]
3. fluoroapatite [ 3Ca3(PO4)2.CaF2 ]
Traces of F2 are present in soil, river water plants and bones and teeth
of animals too.
3. Sea water contains chlorides, bromides and iodides of
Na , K, Mg, Ca but is mainly NaCl solution (2.5% by mass) .
4. Along with NaCl , the deposits of dried up seas contains -
1. Carnallite, KCl.MgCl2.6H2O.
5. Certain forms of marine life contain I2 in their systems;
Ex. - Various seaweeds, contain upto 0.5% of I2
Chile saltpetre contains upto 0.2% of sodium iodate.

Variation of periodic properties
1. Electronic configuration -
1. The valence shell electronic configuration of these electrons is
ns2np5.
2. Thus, there are 7 electrons in the outermost shell of these
elements.
3. The element misses out on the octet configuration by one
electron.
4. Thus, these elements look out to either lose one electron and
form a covalent bond or gain one electron and form an ionic
bond.
5. Therefore, these are very reactive non-metals.

Variation in periodic properties -
2. Atomic and Ionic Radii –
1. The halogens have the smallest atomic radii in their respective
periods due to maximum effective nuclear charge.
2. The atomic radius of fluorine like the other elements of second
period is extremely small.
3. Atomic and ionic radii increase from fluorine to iodine
due to increasing number of quantum shells.
3. Ionisation Enthalpy –
1. They have little tendency to lose electron. Thus they have very high
ionisation enthalpy.
2. Due to increase in atomic size, ionisation enthalpy
decreases down the group.

Variation in periodic properties -
4. Electron Gain Enthalpy –
1. Halogens have maximum negative electron gain enthalpy
in the corresponding periods.
2. This is due to the fact that the atoms of these elements have only one
electron less than stable noble gas configurations.
3. Electronegativity - They have very high electronegativity. The
electronegativity decreases electronegativity down the
group. Fluorine is the most electronegative element in
the periodic table.
4. X-X Bond dissociation enthalpies-
* Expected trend: Cl – Cl > Br – Br > I – I.
Reason - For this anomaly the reason is that the relatively large
electron-electron repulsion among the lone pairs in F2 molecule where
they are much closer to each other than in case of Cl2.

Q. Although electron gain enthalpy of fluorine is less negative as
compared to chlorine, fluorine is a stronger oxidising agent than
chlorine. Why?
5. Electron gain enthalpy -
The electron gain enthalpies of the elements of the group
becomes less negative down the group.
Reason- The negative electron gain enthalpy of fluorine is
less than that of chlorine. It is due to small size of fluorine
atom. As a result, there are strong interelectronic repulsions
in the relatively small 2p orbitals of fluorine and thus, the
incoming electron does not experience much attraction.
Ans. It is due to-
(i) Low enthalpy of dissociation of F-F bond .
(ii) High hydration enthalpy of F– .

Period 12
Learning objective-
1. Physical properties of Cl2.
2. Chemical properties of Cl2 .
3. uses of Cl2.

Physical properties
Physical properties :
The group 17 elements are found in diverse physical states. For
example, Fluorine and Chlorine are gases. On the other hand,
Bromine is a liquid and Iodine is solid.
Colour: These elements have a variety of colours. For example,
while Fluorine is pale yellow in colour, Iodine is dark violet in
colour.
Solubility: Florine and Chlorine are soluble in water. On the other
hand, Bromine and Iodine are very less soluble in water.
Melting and boiling points: Melting and boiling points of
these elements increase as we move down the group from Fluorine
to Iodine. Thus, Fluorine has the lowest boiling and melting points.

Chemical Properties
1. Oxidation states and trends in chemical reactivity –
1. All the halogens exhibit –1 oxidation state.
2. But Cl2, Br2, I2 exhibit + 1, + 3, + 5 and + 7 oxidation states
also as explained below:
1
.

Continued.------
The higher oxidation states of Cl2, Br2, I2 are seen mainly
when the halogens combine with the small and highly
electronegative F2 and O2 atoms.
e.g., in interhalogens, oxides and oxoacids.
The oxidation states of +4 and +6 occur in these oxides and
oxoacids of chlorine and bromine. The fluorine atom has no d
orbitals in its valence shell and therefore cannot expand its
octet. Being the most electronegative, it exhibits only –1
oxidation state.
2. Reactivity of halogens-
All the halogens are highly reactive. They react with metals
and non-metals to form halides.
The reactivity of the halogens decreases down the group.
Oxidising power- The ready acceptance of an electron is the
reason for the strong oxidising nature of halogens. F2 is the
strongest oxidising halogen and it oxidises other halide ions
in solution or even in the solid phase.

2. Reactivity of halogens- continued------
For example- A halogen oxidises halide ions of higher
atomic number.
F2 + 2X
–
→ 2 F
–
+ X2 (X = Cl, Br or I)
Cl2 + 2X
–
→ 2 Cl
–
+ X2 (X = Br or I)
Br2 + 2I– → 2 Br – + I2
The oxidising power decreases down the group which can
be confirmed by their standard electrode potentials .
The relative oxidizing power of halogens can further be shown by
their reactions with water.
2F2 (g) + 2H2O (l) → 4H+ (aq) + 4F− (aq) + O2 (g)
X2 (g) + H2O ( l ) → HX (aq) + HOX (aq)
(where X = Cl or Br)
4I− ( aq) + 4H+ ( aq) + O2 (g) → 2I2 (s) + 2H2O (l)
Most of the reactions of F2 are exothermic (due to the small
and strong bond formed by it with other elements

Do it yourself
Q.1.Considering the parameters such as bond
dissociation enthalpy, electron gain enthalpy
and hydration enthalpy, compare the
oxidising power of F2 and Cl2.
Q.2 Give two examples to show the anomalous
behaviour of fluorine.
Q.3 .Sea is the greatest source of some
halogens. Comment.

3.Reactivity towards hydrogen:
Reactivity towards hydrogen:
Halogen + Hydrogen  Halide
Note- Affinity for hydrogen decreases from F2 to I2.
They dissolve in water to form hydrohalic acids ( table 7.9)
The acidic strength of these acids varies in the order:
HF < HCl < HBr < HI. ( Order of strength)
The stability of these halides decreases down the group –
H–F > H–Cl > H–Br > H–I. ( Order of stability)
Reason- due to decrease in bond (H–X) dissociation enthalpy

4. Reactivity towards oxygen
Halogens form many oxides with O2 but most of them are unstable.
1. F2- a. These oxides are essentially oxygen fluorides because of the
higher electronegativity of fluorine than oxygen.
b. O2F2 oxidises plutonium to PuF6 and the reaction is used in
removing plutonium as PuF6 from spent
2.Cl2 , Br2, I2 – a.Form oxides with oxidation states from +1 to +7.
b. The order of stability of oxides formed is I > Cl > Br.
Reason- Due to kinetic and thermodynamic factors . The higher
oxides of halogens tend to be more stable than the lower ones.
3. Cl2 –Forms Cl2O, ClO2, Cl2O6 and Cl2O7 which are highly
reactive oxidising agents and tend to explode. ClO2 is used as a
bleaching agent for paper pulp and textiles industry.
4. Br2 - The bromine oxides, Br2O, BrO2 , BrO3 are the least stable
halogen oxides (middle row anomally) and exist only at low
temperatures. They are very powerful oxidising agents.

5. Iodine- - The iodine oxides, I2O4 , I2O5, I2O7 are insoluble solids
and decompose on heating. I2O5 is a very good oxidizing agent and is
used in the estimation of carbon monoxide.
5. Reactivity with metals: Halogens + Metals = metal halides.
The ionic character of the halides decreases in the order
MF > MCl > MBr > MI where M is a monovalent metal.
Note- If a metal exhibits more than one oxidation state, the halides in
higher oxidation state will be more covalent than the one in lower
oxidation state .
e.g., SnCl4, PbCl4, SbCl5 and UF6 are more covalent than SnCl2, PbCl2,
SbCl3 and UF4 respectively.

Solved question
Q.1. Fluorine exhibits only –1 oxidation state whereas other
halogens exhibit + 1, + 3, + 5 and + 7 oxidation states also.
Explain.
ANS. Fluorine is the most electronegative element and cannot exhibit
any positive oxidation state. Other halogens have d orbitals and
therefore, can expand their octets and show + 1, + 3, + 5 and + 7
oxidation states .
Intext question ( for practice) -
Q.1.Considering the parameters such as bond dissociation
enthalpy, electron gain enthalpy and hydration enthalpy,
compare the oxidising power of F2 and Cl2.
Q.2. Give two examples to show the anomalous behaviour of
fluorine.
Q.3. Sea is the greatest source of some halogens. Comment.

Chlorine (Cl2) –
Discovered in 1774 by Scheele by the action of
HCl on MnO2.
1. Preparation-It can be prepared by one of the following
methods.
(i)By heating manganese dioxide with concentrated
hydrochloric acid.
MnO2 + 4HCl → MnCl2 + Cl2 + 2H2O
However, a mixture of common salt and concentrated
H2SO4 is used in place of HCl.
4NaCl + MnO2 + 4H2SO4 → MnCl2 + 4NaHSO4 +
2H2O + Cl2
(ii)By the action of HCl on potassium permanganate.
2KMnO4 + 16HCl → 2KCl + 2MnCl2 + 8H2O + 5Cl2

2. Manufacture of chlorine
(i)Deacon’s process: By oxidation of HCl gas by
atmospheric oxygen in the presence of CuCl2
(catalyst) at 723 K.
4HCl + O 2 CuCl2 → 2Cl 2 + 2H2O
(ii)Electrolytic process: Cl2 is obtained by the
electrolysis of brine (conc. NaCl solution). Cl2 is
liberated at anode. It is also obtained as a by–
product in many chemical industries.

3. Properties-
1. It’s is a pungent greenish yellow gas with suffocating odour.
2. It is about 2-5 times heavier than air.
3. It can be liquefied easily into greenish yellow liquid which
boils at 239 K.
4. It is soluble in water.
Chemical properties -
1. Cl2 reacts with a number of metals and non-metals to
form chlorides.
2Al + 3Cl2 → 2AlCl3 ; P4 + 6Cl2 → 4PCl3
2Na + Cl2 → 2NaCl; S8 + 4Cl2 → 4S2Cl2
2Fe + 3Cl2 → 2FeCl3 ;
2. It has great affinity for hydrogen. It reacts with
compounds containing hydrogen to form HCl.

Continued- ----
H2 + Cl2 → 2HCl H2 S + Cl2 → 2HCl + S
C10 H16 + 8Cl2 → 16HCl + 10C
3. With excess NH3, Cl2 gives N2 and NH4Cl .
8NH3 + 3Cl2 → 6NH4Cl + N2 ; ( ammonia in excess)
with excess Cl2 , NCl3 (explosive nitrogen trichloride ) is formed.
NH3 + 3Cl2 → NCl3 + 3HCl (chlorine in excess)
4. Reaction with alkali –
1. Alkali (cold/dil.) + Cl2 → Mix. of chloride and hypochloride.
2NaOH + Cl2 → NaCl + NaOCl + H2O
2. Alkali (hot / conc ) + Cl2 → Mix. of chloride and Chlorate
6 NaOH + 3Cl2 → 5NaCl + NaClO3 + 3H2O
3. With dry slaked lime it gives bleaching powder.
2Ca(OH)2 + 2Cl2 → Ca(OCl)2 + CaCl2 + 2H2O
* Bleaching powder

Continued.
Bleaching powder-
[ composition - Ca(OCl)2.CaCl2.Ca(OH)2.2H2O ]
5. Reaction with hydrocarbons –
1. Cl2 + saturated hydrocarbons → substitution products
CH4 + Cl2 → CH3Cl + HCl
2. Cl2 + unsaturated hydrocarbons → addition products
C2H4 + Cl2 → C2H4
Ethene 1,2-Dichloroethan
6. Bleaching properties- Cl2 water , if kept still , loses its
yellow colour due to the formation of HCl and HOCl.
Hypochlorous acid (HOCl) so formed, gives nascent O2
which adds for oxidising and bleaching properties of Cl2.
continued ……..
*

1. It oxidises ferrous to ferric, sulphite to sulphate,
sulphur dioxide to sulphuric acid and iodine to iodic
acid.
2FeSO4 + H2SO4 + Cl2 → Fe2(SO4)3 + 2HCl
Na2SO3 + Cl2 + H2O → Na2SO4 + 2HCl
SO2 + 2H2O + Cl2 → H2SO4 + 2HCl
I2 + 6H2O + 5Cl2 → 2HIO3 + 10HCl
(ii) A powerful bleaching agent ; ( due to oxidation )
Cl2 + H2O → 2HCl + O
Coloured substance + O → Colourless substance
 It bleaches vegetable or organic matter in the
presence of moisture.
Bleaching effect of chlorine is permanent.

a. For bleaching woodpulp (required for the manufacture of
paper and rayon) and bleaching cotton and textiles,
b. In the extraction of gold and platinum
c. In the manufacture of dyes, drugs and organic
compounds such as CCl4, CHCl3, DDT, refrigerants, etc.
d. In sterilizing drinking water .
e. In the preparation of poisonous gases like-
1. Phosgene (COCl2),
2. Tear gas (CCl3NO2),
3. Mustard gas (ClCH2CH2SCH2CH2Cl).

Assignments no 11
Q.1 Write the balanced chemical equation for the reaction of
Cl2 with hot and concentrated NaOH. Is this reaction a
disproportionation reaction? Justify.
Q.2. Give the reason for bleaching action of Cl2.
Q.3. Name two poisonous gases which can be prepared from
chlorine gas.

Period 13
Learning objective-
HYDROGEN CHLORIDE
 Preparation
Properties
Uses

Hydrogen chloride (HCl)
Discovered by –Glauber
( in 1648 by heating NaCl with conc. H2SO4 )
1. Lab. Preparation
a. Common salt + Conc. H2SO4 → heated to get HCl
NaCl + H2SO4
420 K NaHSO4 + HCl
b. NaHSO4 + NaCl 823 K Na2SO4 + HCl
*HCl gas can be dried by passing through conc, H2SO4.
*2. Physical Properties –
1.It is a colourless and pungent smelling gas.
2.It is easily liquefied to a colourless liquid (b.p.189 K)
and freezes to a white crystalline solid (f.p. 159 K).
3. It is extremely soluble in water and ionises as below:
HCl (g) + H2 O ( l ) → H3 O+ (aq) + Cl− (aq) Ka = 107

Chemical Properties
1. Its aqueous solution is called hydrochloric acid. High value of Ka
indicates that it is a strong acid in water.
2. It reacts with NH3 and gives white fumes of NH4Cl.
NH3 + HCl → NH4Cl
3. Aquaragia-
Mix. of 3 parts of conc. HCl + 1 part of concentrated HNO3
Used for dissolving noble metals, e.g., gold, platinum.
Au + 4H+ + NO3
− + 4Cl− → AuCl−
4 + NO + 2H2O
3Pt + 16H+ + 4NO3
− + 18Cl− → 3PtCl2
6
− + 4NO + 8H2O
4. Hydrochloric acid decomposes salts of weaker acids, e.g.,
carbonates, hydrogencarbonates, sulphites, etc.
Na2CO3 + 2HCl → 2NaCl + H2O + CO2
NaHCO3 + HCl → NaCl + H2O + CO2
Na2SO3 + 2HCl → 2NaCl + H2O + SO2

Uses : It is used
(i) in the manufacture of Cl2, NH4Cl and glucose (from corn starch),
(ii) for extracting glue from bones and purifying bone black, (iii) in
medicine and as a laboratory reagent.
Example of problem-
Q. When HCl reacts with finely powdered iron, it forms ferrous
chloride and not ferric chloride. Why?
Ans. It’s reaction with iron produces H2.
Fe + 2HCl → FeCl 2 + H2
Liberation of H2 prevents the formation of ferric chloride.

Interhalogen Compounds
*When two different halogens react with each other,
interhalogen compounds are formed.
* They can be assigned general compositions as XX′ , XX3′,
XX5′ and XX7′ where X is halogen of larger size and X′ of smaller
size and X is more electropositive than X′.
* As the ratio between radii of X and X′ increases, the number
of atoms per molecule also increases. Thus, iodine (VII)
fluoride should have maximum number of atoms as the ratio
of radii between I and F should be maximum. That is why its
formula is IF7 (having maximum number of atoms).

Preparation
Preparation
By the direct combination or by the action of halogen on lower
interhalogen compounds.
1. Cl2 + F2 437 K 2ClF ;
(equal volume)
2. I2 + 3Cl2 → 2ICl3
3. Cl2 + 3F2→ 2ClF3 ( at 573K)
4. Br2 + 3F2
5. I2 + Cl2 → 2ICl;
6. Br2 + 5F2 → 2BrF5

Some Properties of Interhalogen Compounds

Solved problem-
Q. Deduce the molecular shape of BrF3 on the basis of VSEPR
theory.
Solution – The central atom Br has 7 electrons in the valence shell.
1. 3 of these will form electron-pair bonds with
3 F atoms leaving behind 4 electrons.
2. Thus, there are 3 bond pairs and 2 lone pairs
which will occupy the corners of a
trigonal bipyramid as per VSEPR theory .
3. The 2 lone pairs will occupy the equatorial positions
to minimise lone pair-lone pair and the bond pair-lone pair repulsions which
are greater than the bond pair-bond pair repulsions.
4. In addition, the axial F atoms will be bent towards the equatorial
F in order to minimise the lone-pair-lone pair repulsions. The shape
would be that of a slightly bent ‘T’.

Period 14
Group 18 Elements ( Noble gases)
Occurrence
Electronic configuration.
Ionization enthalpies
Atomic radius
Electron gain enthalpies

1.Occurance -
1. All the noble gases except radon occur in the atmosphere.
2. Argon is about 1% by volume in dry air in atmosphere.
3. Helium and Neon are found in minerals of radioactive
origin e.g., pitchblende, monazite, cleveite.
4. Helium - The main commercial source is natural gas.
5. Xe and Rn - Radon is obtained as a decay product of
226Ra.
226
88
Ra →
222
86
Rn +
4
2
He

1. Electronic configuration -
All noble gases have general electronic configuration ns2np6
except helium which has 1s2 . Many of the properties of
noble gases including their inactive nature are due to their closed
shell structures.
2. Ionization enthalpy- Due to stable electronic configuration
these gases exhibit very high ionization enthalpy. It decreases
down the group with increase in atomic size.
3. Atomic Radius- Atomic radii increase down the group with
increase in atomic number.
4. Electron gain enthalpy - Since noble gases have stable
electronic configurations, they have no tendency to accept the
electron and therefore, have large positive values of electron gain
enthalpy.

Chemical reactivity
Chemical Properties
1. In general, noble gases are least reactive.
2. Reasons for inertness-
a. The noble gases except helium (1s2) have completely filled
ns2np6 electronic configuration in their valence shell.
b. They have high ionization enthalpy and more positive
electron gain enthalpy.
Few important compounds of noble gasses –
1.KrF2 – Kripton difluoride.
2.RnF2 – Radon difluoride , its only identifies but not
isolated.
3. Ar ,Ne and helium yet not identified to make
compounds.

Xenon-oxygen compounds
Hydrolysis of XeF4 and XeF6 with water gives Xe03 -
6XeF4 + 12 H2O → 4Xe + 2Xe03 + 24 HF + 3 O2
XeF6 + 3 H2O → XeO3 + 6 HF
Partial hydrolysis of XeF6 gives
Oxyfluorides, XeOF4 and XeO2F2.
XeF6 + H2O → XeOF4 + 2 HF
XeF6 + 2 H2O → XeO2F2 + 4HF
• XeO3 is a colourless explosive solid and
•Has a pyramidal molecular structure (Fig.
•XeOF4 is a colourless volatile liquid
•Has a square pyramidal molecular

Florides of Xenon
Pyramidal structure of XeO3

He –
1. Filling balloons for meteorological observations.
2. It is also used in gas-cooled nuclear reactors.
3. Liquid He is used as cryogenic agent .
Ne - Used in making lamps/ fluorescent tubes and for
display purpose.
Ar. – Mainly used for giving inert atmosphere in high
temprature metallurgical purpose.
There are no significant uses of Xenon and
Krypton. They are used in light bulbs designed
for special purposes.

Ppt xii p block elements , By Dr Mona Srivastava

Ppt xii p block elements , By Dr Mona Srivastava

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