P block elements

P block elements
Mr. Vijaykumar Nazare

The p-Block Elements
p-Block Elements (Group 13 to 18 )
s-Block Elements (group 1 and 2 )
s and p- block
Representative Elements or Main Group
Elements.
27-Oct-19 3Vijaykumar Nazare

Outermost electronic configuration varies from
ns2np1 to ns2np6
• Group 13 elements--- ns2np1
• Group 14 elements--- ns2np2

The p-Block
Elements
 Most of p-Block elements are non-metals.
 They have variable oxidation states.
 They form acidic oxides
 They impart no characteristic colour to the
flame
 Generally they form covalent compounds.
Halogens form salts with alkali metals
Main points (properties)

• They have high ionization potentials.
• They have very large electron gain
enthalpies.
• They are solids/liquids/gases at room
temperature (Br is liquid)
• The aqueous solutions their oxides are
acidic in nature.
Main points (properties) cont…d

THE p -BLOCK ELEMENTS
• why p-block elements consist of
only six groups?
• The number of p orbitals is three therefore,
the maximum number of electrons that can
be accommodated in a set of p orbitals is six.
• Therefore, there are six groups of p–block
elements in the periodic table numbering
from 13 to 18.

• Boron, carbon, nitrogen, oxygen, fluorine and
helium head the groups.
• Their valence shell electronic configuration is
ns2np1-6(except for He).

The important oxidation states
exhibited by p-block elements are shown in Table

PROPERTY N P AS SB BI
Atomic
number 7 15 33 51 83
Atomic mass/g mol-1 14.01 30.97 74.92 121.75 208.98
Electronic
configuration [He]2S22p3 [Ne]3S22p3 [Ar]3d104s24p3 [Kr]4d105s25p3 [Xe]4f145d106s26p3
Ionisation I
enthalpy II
(ΔiH/(kj III mol-1)
1402
2856
4577
1012
1903
2910
947
1798
2736
834
1595
1610
703
1610
2466
Electronegetivity 3.0 2.1 2.0 1.9 1.9
Covalent radius/pma 70 110 121 141 148
Ionic radius/pm 171b 212b 222b 76c 103c
Melting point/K 63* 317d 1089e 904 544
Boiling point/K 77.2* 554d 888f 1860 1837
Density/[g cm-3(298
K)] 0.879g 1.823 5.778h 6.697 9.808
Table 7.1: Atomic and Physical Properties of Group 15 Elements
27-Oct-19 Vijaykumar Nazare 10

Group 15 elements
Electronic Configuration
• The valence shell electronic configuration
ns2np3.
• The s orbital is completely filled and p orbitals
are half-filled, making their electronic
configuration extra stable.
• +3 and +5 oxidation state .

7.1.3 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.
• As to Bi only a small increase in covalent
radius is observed due to the presence of
completely filled d orbitals and/or f orbitals in
heavier members.

7.1.4 Ionisation Enthalpy
• Ionisation enthalpy : decreases down the group
due to gradual increase in atomic size.
• Ionisation enthalpy of group 15 elements greater
than group 14 elements:Because of the extra
stable half-filled p orbitals electronic
configuration and smaller size of group 15
elements .
• Increase in magnitude of effective nuclear charge.

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

7.1.6 Physical Properties
• Polyatomic nature : Dinitrogen diatomic gas while all
others are solids (Polyatomic P4 ).
• Metallic character : increases down the group.
Nitrogen and phosphorus are non-metals,
arsenic and antimony metalloids
bismuth is a metal.
This is due to decrease in ionisation enthalpy and increase
in atomic size.
• The boiling points, increase from top to bottom in
the group .
• The melting point increases upto arsenic and then
decreases upto bismuth.

Anomalous properties of nitrogen
Anamalous property is due to
1) its small size
2) High electronegativity
3) High ionisation enthalpy
4) Non-availability of d orbitals.

Anomalous property
• Nitrogen form p π -p π multiple bonds .
• Bond enthalpy (941.4 kJ mol–1 ) is very high.
• Heavier elements do not form p π -pπ bonds as
their atomic orbitals are so large and diffuse that
they cannot have effective overlapping.
• phosphorus (P–P), arsenic(As–As) and
antimony (Sb–Sb)form single bonds and bismuth
forms metallic bonds in elemental state.

Anamalous property
• N-N single bond is weaker than P-P single bond .
Because
• Bond length is short in N-N .
• High interelectronic repulsion of the non-bonding
electrons.
• Therefore catenation is weak in nitrogen .

• Nitrogen cannot form bond with transition
elements :
• Absence of d orbitals in its valence shell.
• Nitrogen cannot form dπ –pπ bond as the heavier
elements can e.g., R3P = O or R3P = CH2 (R = alkyl
group).
• Phosphorus and arsenic can form dπ –dπ bond
with transition metals when their compounds like
P(C2H5)3 and As(C6H5)3 act as ligands.

(i) Reactivity towards hydrogen
• Hydride formation : Group 15 form EH3
( E = N, P, As, Sb or Bi )
• The hydrides show regular gradation in their
properties.
• Stability : decreases from NH3 to BiH3 due to increase
in atomic size , decrease in bond dissociation enthalpy.
• Reducing character :Increases ,due to small bond
dissociation enthalpy ,covalent character decreases.
• Basicity :decreases NH3 > PH3 > AsH3 > SbH3 > BiH3.
• NH3 Is strong base :Small size and high electron
density, has lone pair .

(ii) Reactivity towards oxygen
• Form two types of oxides: E2O3 and E2O5.
• The oxide in higher oxidation state is more acidic.
Their acidic character decreases down the group.
• Nitrogen atom has small atomic size ,strong pull of
electron pair between O-H bond ,releases the H+
ion .
• Effect decreases as atomic size increases .
• The oxides of type E2O3 of nitrogen and
phosphorus are purely acidic, arsenic and antimony
amphoteric ,bismuth predominantly basic.

(iii) Reactivity towards halogens
• Halides formation: MX3 and MX5.
• Nitrogen does not form pentahalide due to non-
availability of the d orbitals in its valence shell,
contains only 1 - s and 3- p orbitals .
• Pentahalides are more covalent than trihalides.
• All the trihalides (covalent nature) of these
elements except those of nitrogen are stable.
• In case of nitrogen, only NF3 is known to be
stable. Trihalides except BiF3 are predominantly
covalent in nature.

Reactivity towards metals
• Form binary compounds (having +3 oxidation)
• Ca3N2 (Calcium Nitride) ,Ca3P2 (Calcium Phosphide)
Na3As2(Sodium Arsenide ), Zn3Sb2(ZincAntimonide)
Mg3Bi2 (Magnesium Bismuthide)

• Though nitrogen exhibits +5 oxidation state, it does
not form pentahalide. Give reason.
• Solution
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.
• PH3 has lower boiling point than NH3. Why?
• Solution
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.

7.2 Dinitrogen
Preparation :( NH4CI, NaNO2 / (NH4)2Cr2O7 / Ba(N3)2 )
Laboratory : dinitrogen is prepared by treating an aqueous solution
of ammonium chloride with sodium nitrite.

NH4CI(aq) + NaNO2(aq) → N2(g) + 2H2O(l) + NaCl (aq)

It can also be obtained by the thermal decomposition of
ammonium dichromate.
 Very pure nitrogen :thermal decomposition of sodium or
barium azide.
Ba(N3)2 → Ba + 3N227-Oct-19 Vijaykumar Nazare 25

Properties
• Dinitrogen : colourless, odourless, tasteless and non-toxic gas.
• Two stable isotopes: 14N and 15N.
• Low Solubility in water , low freezing and boiling points
• Inert at room temperature because of the high bond enthalpy
of N ≡N bond. Reactivity increases rapidly with rise in
temperature.
• At higher temperatures, form ionic nitrides and with non-
metals, covalent nitrides.
• Reaction with hydrogen to form ammonia:
• Reaction with dioxygen (at about 2000 K) form nitric oxide,
NO.

Uses:
• Manufacture of ammonia and industrial
chemicals containing nitrogen, (e.g., calcium
cyanamide).
• Finds use where an inert atmosphere is
required (e.g., in iron and steel industry, inert
diluent for reactive chemicals).
• Refrigerant to preserve biological materials,
food items and in cryosurgery etc.

Ammonia Preparation
• Ammonia :present in air and soil formed by decay of
nitrogenous organic matter e.g., urea.
NH2CONH2 + 2H2O → ( NH4 )2CO3 → 2NH3 + H2O + CO2
• On a small scale ammonia is obtained from ammonium salts
which decompose when treated with caustic soda or lime.
2NH4Cl + Ca(OH)2 → 2NH3 + 2H2O + CaCl2
(NH4)2 SO4 + 2NaOH → 2NH3 + 2H2O + Na2SO4
• On a large scale, ammonia is manufactured by Haber’s process.
• N2(g) + 3H2(g) → 2NH3(g) ΔfH° = – 46.1 kJ mol−1

Properties of Ammonia
• colourless gas with a pungent odour
• In the solid and liquid states, it is associated through
hydrogen bonds , high melting and boiling points .
• Structure : trigonal pyramidal. It has three bond pairs and
one lone pair of electrons .
• Ammonia gas : 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)
• Lewis base :The presence of lone pair of electrons on
nitrogen atom. donates the electron pair and forms complex
with metal ions .
• Applications :detection of metal ions such as Cu2+ , Ag+

Applications
• Cu2+(aq) + 4NH3(aq) → [Cu(NH3)4]2+ (aq)
(blue) (deep blue)
• Ag+(aq) + Cl− ( aq ) → AgCl ( s )
(colourless) (white ppt)
• AgCl ( s ) + 2NH3 ( aq ) → Ag ( NH3 )2Cl ( aq )
(white ppt) (colourless)

Uses:
• Produce nitrogenous fertilisers (ammonium
nitrate, urea, ammonium phosphate and
ammonium sulphate)
• Manufacture of some inorganic nitrogen
compounds, eg. nitric acid.
• Liquid ammonia is also used as a refrigerant.

Oxides of Nitrogen

Nitric Acid
• Nitrogen forms oxoacids : H2N2O2 (hyponitrous acid)
• HNO2 (nitrous acid) and HNO3 (nitric acid). HNO3 is the most
important.
• Preparation:
• Laboratory :heating KNO3 or NaNO3 and concentrated H2SO4 in a
glass retort.
NaNO3 + H2SO4 → NaHSO4 + HNO3
• Large scale : Ostwald’s process. catalytic oxidation of NH3 by
atmospheric oxygen.
• Nitric oxide thus formed combines with oxygen giving NO2.
2NO ( g ) + O2 ( g )→2NO2 ( g )
• Nitrogen dioxide so formed, dissolves in water to give HNO3.
3NO2 ( g ) + H2O ( l ) → 2HNO3 ( aq ) + NO ( g )

Properties
• It is a colourless liquid,
• Structure : planar molecule .
Nitric acid behaves as a strong acid giving hydronium and nitrate ions.
HNO3(aq) + H2O(l) → H3O+ (aq) + NO3
- (aq)
• Strong oxidising agent and attacks most metals except noble metals such
as gold and platinum.
• 3Cu + 8 HNO3(dilute) → 3Cu(NO3)2 + 2NO + 4H2O
Cu + 4HNO3(conc.) → Cu(NO3)2 + 2NO2 + 2H2O
• Zinc reacts with dilute nitric acid to give N2O and with concentrated acid
to give NO2.
4Zn + 10HNO3(dilute) → 4 Zn (NO3)2 + 5H2O + N2O
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.

Properties
• Reaction with Non- metals: Iodine is oxidised
to iodic acid, carbon to carbon dioxide,
sulphur to H2SO4, and phosphorus to
phosphoric acid.
• I2 + 10HNO3 → 2HIO3 + 10NO2 + 4H2O
C + 4HNO3 → CO2 + 2H2O + 4NO2
S8 + 48HNO3 → 8H2SO4 + 48NO2 + 16H2O
P4 + 20HNO3 → 4H3PO4 + 20NO2 + 4H2O

Uses:
• Manufacture of ammonium nitrate for fertilisers
and other nitrates for use in explosives and
pyrotechnics.
• Preparation of organic nitro compounds
nitroglycerin, trinitrotoluene and other.
• Other major uses are in the pickling of stainless
steel, etching of metals and an oxidiser in rocket
fuels.

Sr.
no
White Phosphorus Red Phosphorus Black phosphorus
1 Transparent , waxy solid
,poisonous .
Odourless ,non-poisonous
white P4 + 573 K → Red P4
iron grey lustre
Red P4 + 803 K → α-black P4
white P4 + 4373 K → β-black P4
Two forms α and β
2 insoluble in water but
soluble in CS2
insoluble in water and CS2 Sublime,opaque,monoclinic
3 Glows in dark Does not Glow in dark Does not glow in dark
4 Burns in air
P4 + 5O2 → P4O10
β-black P4 Does not burn in air
5 Less stable ,reactive due
to angular strain
Less reactive Less reactive
6 tetrahedral P4 molecule polymeric, chains of P4
tetrahedra linked together .
Layered structure

Phosphine PH3
• Preparation : Reaction of calcium phosphide with water or
dilute HCl.
Ca3P2 + 6H2O → 3Ca(OH)2 + 2PH3
Ca3P2 + 6HCl → 3CaCl2 + 2PH3
• Laboratory : Heating white phosphorus with concentrated
NaOH solution in an inert atmosphere of CO2.
P4 + 3NaOH + 3H2O → PH3 + 3NaH2 PO2
(sodium hypophosphite)
When pure, it is non inflammable but becomes
inflammable owing to the presence of P2H4 or P4 vapours.
• To purify
PH4I + KOH → KI + H2O + PH3

Properties
• Colourless gas with rotten fish smell and highly
poisonous.
• It explodes in contact with oxidising agents like HNO3,
Cl2 and Br2 vapours.
• Slightly soluble in water. The solution of PH3 in water
decomposes in presence of light giving red phosphorus
and H2.
• When absorbed in copper sulphate or mercuric
chloride solution, phosphides are obtained.
3CuSO4 + 2PH3 → Cu3 P2 + 3H2SO4
3HgCl2 + 2PH3 → Hg3P2 + 6HCl
• Phosphine is weakly basic and like ammonia, gives
phosphonium compounds with acids e.g.,
PH3 + HBr → PH4 Br

Uses:
• The spontaneous combustion of phosphine is
technically used in Holme’s signals.
• Containers containing calcium carbide and
calcium phosphide are pierced and thrown in
the sea when the gases evolved burn and
serve as a signal.
• It is also used in smoke screens.

Phosphorus Halides
Two types of halides, PX3 (X = F, Cl, Br, I) and PX5 (X = F, Cl, Br).
• Phosphorus Trichloride
• Preparation :
• The reaction of white phosphorus
with excess of dry chlorine.
P4 + 6Cl2 → 4PCl3
• Action of thionyl chloride with
white phosphorus.
P4 + 8SOCl2 → 4PCl3 + 4SO2 + 2S2Cl2
• Phosphorus Pentachloride
• Preparation :
• The reaction of white
phosphorus with excess of dry
chlorine.
P4 + 10Cl2 → 4PCl5
• Action of SO2Cl2 on
phosphorus.
P4 + 10SO2Cl2 → 4PCl5 + 10SO2

Properties
• Phosphorus Trichloride
1. Colourless oily liquid
2. hydrolysed in presence of
moisture.
PCl3 + 3H2O →H3PO3 + 3HCl
3. Reacts with organic compounds
containing –OH group.
3CH3COOH + PCl3 → 3CH3COCl
+H3PO3
3C2H5OH + PCl3 → 3C2H5Cl + H3PO3
4. It has a pyramidal shape sp3 .
Phosphorus Pentachloride
1. Yellowish white powder
2. Hydrolyses in presence of
moisture.
PCl5 + H2O → POCl3 + 2HCl
POCl3 + 3H2O → H3PO4 +
3HCl
3. Reacts with organic
compounds containing –OH
group .
C2H5OH + PCl5 → C2H5Cl + POCl3 + HCl
CH3COOH + PCl5 → CH3COCl + POCl3
+HCl
5. Trigonal bipyramidal sp3d

• Metals on heating with PCl5 .
2Ag + PCl5 → 2AgCl + PCl3
Sn + 2PCl5 → SnCl4 + 2PCl3
• Trigonal bipyramidal structure .Two axial bonds
are longer than equatorial bonds. This is due to
the fact that the axial bond pairs suffer more
repulsion as compared to equatorial bond pairs.
• [PCl4]+ is tetrahedral and the anion, [PCl6]-
octahedral

NAME FORMULA
OXIDATION STATE OF
PHOSPHORUS
CHARACTERISTIC
S BONDS AND
THEIR NUMBER PREPARATION
Hypophosphorus
(phosphinic) H3PO2 +1
One P – OH
Two P – OH
One P = O white P4 + alkali
Orthophosphorous
(phosphonic) H3PO3 +3
Two P – OH
One P – OH
One P = O P2O3
Pyrophosphorous H4P2O5 +3
Two P – OH
two P – OH
Two P = O PCl3 + H3PO3
Hypophosphoric H4P2O6 +4
Four P – OH
two P – OH
One P = O red P4 + alkali
Orthophosphoric H3PO4 +5
Three P – OH
One P – OH P4O10+H2O
Pyrophosphoric H4P2O7 +5
Two P – OH
Two P – OH
One P-O-P heat phosphoric acid
Metaphosphoric (HPO3)n +5
Three P – OH
Three P – OH
Three P-O-P
phosphorus acid +
Br2, heat in a sealed
tube
Table 7.5: Oxoacids of Phosphorus

Oxoacids of Phosphorus

• Acids having P–H bond have strong reducing
properties. Thus, hypophosphorous acid H3PO2 is a
good reducing agent as it contains two P–H bonds
and reduces.
4 AgNO3 + 2H2O + H3PO2 → 4Ag + 4HNO3 + H3PO4
• These P–H bonds are not ionisable to give H+ and
do not play any role in basicity.
• H atoms which attached with oxygen in P–OH
form are ionisable and cause the basicity. Thus,
H3PO3 and H3PO4 are dibasic and tribasic,
respectively as the structure of H3PO3 has two P–
OH bonds and H3PO4 three.

Group 16 Elements
• Oxygen, sulphur, selenium, tellurium and
polonium(radioactive).(chalcogens – ore forming)
• Derived from Greek word for brass and points to
the association of sulphur and its congeners with
copper.
• Copper minerals contain oxygen or sulphur and
other members of the group.
• Present I earth crust ,gypsum,epsum,pyrite,zinc
blend ,H2S in volcanoes,protein,garlic,onion,hair .

1. Electronic Configuration : ns2 np4
2. Atomic and Ionic Radii : Increases
Due to increase in the number of shells
3. Ionisation Enthalpy : Decreases
due to increase in size .
Grop16 has lower I.E than Group15 .
due to the fact that Group 15 elements have extra stable
half- filled p orbitals electronic configurations.

1. Electron Gain Enthalpy : Because
of the compact nature of oxygen atom (small size)
e-e repulsion, it has less negative electron gain
enthalpy than sulphur. However, from sulphur
onwards the value again becomes less negative
upto polonium due to increase in size.
2. Electronegativity : F >O >N
electronegativity decreases with an increase in atomic
number or size. Metallic character increases from oxygen
to polonium.

PEOPERTY O S SE TE PO
Atomic number 8 16 34 52 84
Atomic mass/g mol-1 16.00 32.06 78.96 127.60 210.00
Electronic
configuration [He]2s22p4 [Ne]3s23p4 [Ar]3d104s24p4 [Kr]4d105s25p4 [Ar]4f145d106s26p4
Covalent radius/(pm)a 66 104 117 137 146
Ionic radius, E2-/pm 140 184 198 221 230b
Electron gain
enthalpy,/ΔegH kJ
mol-1 -141 -200 -195 -190 -174
Ionisation enthalpy
(ΔiHi)/kJ mol-1 1314 1000 941 869 813
Electronegetivity 3.50 2.44 2.48 2.01 1.76
Density /g cm-3(298
K) 1.32c 2.06d 4.19e 6.25 -
Melting point/K 55 393f 490 725 520
Boiling point/K 90 718 958 1260 1235
Oxidation states -2,-1,1,2 -2,2,4,6 -2,2,4,6 -2,2,4,6 2,4
Table 7.6: Some Physical Properties of Group 16 Elements

Physical Properties
• Radioactive
• Exhibit allotropy
• Melting and boiling point increases due to
increase in atomic mass .

Chemical Properties
1. Oxidation state : -2 ,-1 ,+2 ,+4 ,+6
2. +2 OF2
3. Oxygen does not show +4 and +6 O.S due to
lack of d-orbitals .
4. Stability of +6 oxidation state in higher
elements due to inner pair effect .

Anomalous behaviour of oxygen
• Small size ,high I.E. and high electronegativity.
• The absence of d orbitals in oxygen limits its
covalency to four ,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.

Reactivity with hydrogen
• Hydrides of the type H2E (E = O, S, Se, Te, Po).
• acidic character: increases from H2O to H2Te.
Due to decrease in bond (H–E) dissociation
enthalpy .
• Thermal stability :decrease
bond (H–E) dissociation enthalpy decreases.
• Reducing property: character increases from H2S
to H2Te. Bond length increases .

Reactivity with oxygen
• Oxides : EO2 and EO3 ( E = S, Se, Te ,Po )
• Ozone (O3) and sulphur dioxide (SO2) and (SO3)
are gases while selenium dioxide (SeO2) is solid.
• Reducing property : of dioxide decreases from SO2
to TeO2 .
• Besides EO2 type, sulphur, selenium and tellurium
also form EO3 type oxides (SO3, SeO3, TeO3). Both
types of oxides are acidic in nature.

Reactivity towards the halogens
• Type : EX6, EX4 and EX2 .
• Stability : decreases in the order F− > Cl− > Br− > I− .
• Hexahalides : hexafluorides are only stable halides.
gaseous in nature, octahedral structure sp3d2. Eg.SF6 .
• Tetrafluorides : SF4 - gas, SeF4 -liquid and TeF4 - solid.
Sp3d hybridisation ,have trigonal bipyramidal ,having
lone pair of electrons at equitorial position.
• All elements except selenium form dichlorides and
dibromides. sp3 hybridisation , tetrahedral structure.

Acidic nature
• H2S is less acidic than H2Te. Why?
Solution
Due to the decrease in bond (E–H)
dissociation enthalpy down the group, acidic
character increases.

Dioxygen
• Preparation :
• Laboratory: heating oxygen containing salts such as
chlorates, nitrates and permanganates.
• (ii) Thermal decomposition :
2Ag2O(s) → 4Ag(s) + O2(g) 2Pb3O4(s) → 6PbO(s) + O2(g)
2HgO(s) → 2Hg(l) + O2(g) 2PbO2(s) → 2PbO(s) + O2(g)
• (iii) Decomposition of H2O2 using manganese dioxide.
2H2O2(aq) → 2H2O(l) + O2(g)
• large scale: Electrolysis of water ,release of hydrogen at the
cathode and oxygen at the anode.
• Industrially :from air by first removing carbon dioxide and
water vapour and then, the remaining gases are liquefied
and fractionally distilled to give dinitrogen and dioxygen.

Properties
• Colourless and odourless gas, soluble in water.
• 3 isotopes: 16O ,17O and 18O.
• Paramanetic
• Dioxygen reacts with metals and non-metals except
some metals ( e.g., Au, Pt) and some noble gases.
• Reactions :
2Ca + O2 → 2CaO C + O2 → CO2
4Al + 3O2 → 2Al2O3 2ZnS + 3O2 → 2ZnO + 2SO2
P4 + 5O2 → P4O10 CH4 + 2O2 → CO2 + 2H2O
• Exothermic reaction ,to initiate the reaction, some
external heating is required as bond dissociation
enthalpy of oxgyen-oxygen double bond is high (493.4
kJ mol–1).27-Oct-19 Vijaykumar Nazare 61

Uses
• Respiration and combustion processes,
• oxy acetylene welding (manufacture of steel)
• Oxygen cylinders in hospitals, high altitude flying
and in mountaineering.
• The combustion of fuels, e.g., hydrazines in liquid
oxygen, provides tremendous thrust in
rockets.(L.O as oxidiser in rocket fuel)

Simple Oxides
• Classification : acidic, basic or amphoteric character.
• An acidic oxide: oxide combines with water give acid.
(e.g., SO2, Cl2O7, CO2, N2O5 ).
SO2 + H2O → H2SO3 (only non-metal oxides are acidic)
• Metals in high oxidation state have acidic character
(e.g.Mn2O7, CrO3, V2O5).
• basic oxides :The oxides which give base with water
• (e.g., Na2O, CaO, BaO)
CaO + H2O → Ca(OH)2 (metallic oxides are basic)
• amphoteric oxides : shows both acidic as well as basic
character. Eg. Al2O3
Al2O3(s)+ 6HCl (aq) + 9H2O ( l ) → 2[ Al(H2O)6]3+(aq)+6Cl−(aq )
Al2O3 ( s ) + 6NaOH ( aq ) + 3H2O ( l ) → 2Na3[Al(OH)6](aq)
• neutral oxides :neither acidic nor basic.
• Eg. CO, NO and N2O.27-Oct-19
Vijaykumar Nazare
63

Ozone
• allotropic form of oxygen.
• Formation : from atmospheric oxygen in the
presence of sunlight. ozone layer protects the
earth’s surface from excessive concentration of
ultraviolet (UV) radiations.
• Preparation :
• Stream of oxygen passed through silent electrical
discharge, conversion of oxygen to ozone (10%)
occurs.
• 3O2 → 2O3 Ozonised oxygen ΔHV (298 K) = +142 kJ mol−1
• Endothermic process .27-Oct-19 Vijaykumar Nazare 64

Properties
1. pale blue gas, dark blue liquid and violet-black solid.
2. characteristic smell small concentrations harmless
,higher concentration headache and nausea.
3. Thermodynamically unstable.
4. Strong oxidising agent .
5. decomposition into oxygen results in liberation of
heat (ΔH is negative) and increase in entropy (ΔS is
positive).
6. large negative Gibbs energy change (ΔG) for its
conversion into oxygen.

1. Nitrogen oxides emitted from supersonic jet
aeroplanes depletes ozone layer .
NO ( g ) + O3 ( g ) → NO2 ( g ) + O2 ( g )
2. Use of freons which are used in aerosol sprays
and as refrigerants depletes ozone.
3. high concentrations of ozone is explosive.

Uses
1. germicide, disinfectant and for sterilising
water.
2. bleaching oils, ivory, flour, starch, etc.
3. oxidising agent in manufacture of potassium
permanganate.

Sulphur – Alltropic Forms
Rhombic sulphur (α-sulphur)
1. yellow in colour, m.p. 385.8 K
and specific gravity 2.06.
2. formed on evaporating
solution of sulphur in CS2.
3. Insoluble in water , soluble in
CS2.
4. stable below 369 K and
transforms into β-sulphur
above this temperature .
Monoclinic sulphur (β-sulphur)
1. Its m.p. is 393 K and specific
gravity 1.98.
2. formed by melting rhombic
sulphur in a dish and cooling.
3. It is soluble in CS2.
4. stable above 369 K and
transforms into α-sulphur
below it.
5. At 369 K both the forms are
stable. This temperature is
called transition
temperature.

• Both have S8 molecule ,
• S8 ring is puckered and has a crown shape.
• In cyclo-S6, the ring adopts the chair form .

Sulphur Dioxide
1. Preparation :
• Sulphur is burnt in air or oxygen:
S(s) + O2(g) → SO2 (g)
• Laboratory: treating sulphite with dilute
sulphuric acid.
SO3 (aq) + 2H (aq) → H2O(l) + SO2 (g)
• Industrially : roasting of sulphide ores.
4FeS2 (s ) + 11O2 ( g ) → 2Fe2O3 ( s ) + 8SO2 ( g )

Properties
1. Colourless gas, pungent smell and highly soluble in water ,
reducing agent.
2. Sulphur dioxide reacting with water, forms sulphurous acid.
SO2(g) + H2O(l) → H2SO3(aq)
3. Sodium hydroxide solution, forming sodium sulphite,reacts with
more sulphur dioxide to form sodium hydrogen sulphite.
4. 2NaOH + SO2 → Na2SO3 + H2O
Na2SO3 + H2O + SO2 → 2NaHSO3
5. Sulphur dioxide reacts with chlorine in presence of charcoal
(catalyst) gives 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)

Uses of SO2
• (i) refining petroleum and sugar
• (ii) bleaching wool and silk and
• (iii) as an anti-chlor, disinfectant and preservative.
• To prepare Sulphuric acid, sodium hydrogen
sulphite and calcium hydrogen sulphite (industrial
chemicals)
• Liquid SO2 used as solvent .

Oxoacid of Sulphur
+4 +6 +7 +6

Sulphuric Acid
• Manufacture : (Contact Process )
• (i) burning of sulphur or sulphide ores in air to generate SO2.
S → SO2
• (ii) conversion of SO2 to SO3 by the reaction with oxygen in
the presence of a catalyst (V2O5)
•
(iii) absorption of SO3 in H2SO4 to give Oleum (H2S2O7).
• exothermic, reversible and the forward reaction
• low temperature and high pressure .
• SO3 + H2SO4 → H2S2O7
(Oleum)

Flow diagram for H2SO4

Properties of H2SO4
1. colourless, dense, oily liquid .
2. The chemical reactions due following
characteristics: (a) low volatility (b) strong acidic
character (c) strong affinity for water (d) ability
to act as an oxidising agent.
3. Ionisation of acid in water.
4. H2SO4(aq) + H2O(l) → H3O+ (aq) + HSO4
− (aq); Ka1
= very large ( Ka1 >10)
HSO4 (aq) + H2O(l) → H3O+ (aq) + SO4
2− (aq) ; Ka2
> = 1.2 × 10−2
5. Greater value of(Ka), the stronger is the acid.27-Oct-19 Vijaykumar Nazare 76

Properties of H2SO4
1. because of low volatility used to manufacture more
volatile acids .
2 MX + H2SO4 → 2HX + M2SO4 (X = F, Cl, NO3)
(M = Metal)
• Strong dehydrating agent.
• Strong oxidising agent.
Cu + 2 H2SO4(conc.) → CuSO4 + SO2 + 2H2O
3S + 2H2SO4(conc.) → 3SO2 + 2H2O
C + 2H2SO4(conc.) → CO2 + 2 SO2 + 2 H2O

Uses of H2SO4
1. fertilisers (ammonium sulphate, superphosphate).
2. (a) petroleum refining.
3. (b) pigments, paints and dyestuff intermediates .
4. (c) detergent industry .
5. (d) metallurgical applications (e.g., cleansing
metals before enameling, electroplating and
galvanising .
6. (e) storage batteries .
7. (f) manufacture of nitrocellulose products .
8. (g) a laboratory reagent.

Group 17 Elements
• Fluorine, chlorine, bromine, iodine and astatine
(radioactive).
• halogens (salt forming or salt producers).
• Highly reactive , non-metallic elements .
• Occurance :
• Fluorine :Fluorspar CaF2 ,Cryolite Na3AlF6
• Cl ,Br ,I :Sea water as salt of Na ,K,Mg ,Ca,

1. Electronic Configuration : (ns2 np5 )
2. Atomic and Ionic Radii :
smallest atomic radii due to maximum effective
nuclear charge.
Atomic and ionic radii increase due to increasing
number of quantum shells.
3. Ionisation Enthalpy :
Little tendency to lose electron due to very high
ionisation enthalpy.
Due to increase in atomic size, ionisation
enthalpy decreases down the group.

1. Electron Gain Enthalpy:
Maximum :only one electron less than stable
noble gas configurations.
2. Negative electron gain enthalpy of fluorine is
less than that of chlorine due to small size of
fluorine atom ,strong interelectronic repulsions in
2p orbitals of fluorine , experience less attraction.
3. Electronegativity :
Very high due to increase nuclear charge.
Decreases down the group due to increase
atomic radia.

PEOPERTY F CL BR I ATA
Atomic number 9 17 35 53 85
Atomic mass/g mol-1 19.00 35.42 79.90 126.90 210
Electronic configuration [He]2s22p5 [Ne]3s23p5 [Ar]3d104s24p5 [Kr]4d105s25p5 [Ar]4f145d106s26p5
Covalent radius/(pm)a 64 99 114 133 -
Ionic radius, X-/pm 133 184 196 220 -
Ionisation enthalpy
(ΔiHi)/kJ mol-1 1680 1256 1142 1008 -
Electron gain enthalpy/kJ
mol-1 -333 -349 -325 -296 -
Electronegetivity 4 3.2 3.0 2.7 2.2
ΔHydH(X-)/kJ mol-1 515 381 347 305 -
F2 CL2 BR2 I2 -
Melting point/K 54.4 172.0 265.8 386.6 -
Boiling point/K 84.9 239.0 332.5 458.2 -
Density/g cm-3 1.5 (85)c 1.66(203)c 3.19(273)c 4.49(293)d -
Distance X- X/pm-3 143 199 228 266 -
Bond dissociation
enthalpy/(kJ mol-1 158.8 242.2 192.8 151.1 -
EV/Ve 2.87 1.36 1.09 0.54 -

Physical Properties
• F,Cl - gases, Br - liquid ,I - solid.
• melting and boiling points increase with atomic number.
• Coloured : Due to absorption of radiations in visible region
which results in the excitation of outer electrons to higher
energy level. By absorbing different quanta of radiation, they
display different colours. Eg. F2 – yellow, Cl2 - greenish
yellow, Br2- red and I2-violet colour.
• Bond dissociation enthalpy : F2 < Cl2 >Br2 > I2
• F2 has smaller bond dissociation enthalpy than Cl2 Due to
1. large electron-electron repulsion among the lone pairs in F2
2. Much closer to each other than Cl2.

Chemical Properties
• Oxidation states ,All the halogens exhibit –1 oxidation state. However,
chlorine, bromine and iodine exhibit + 1, + 3, + 5 and + 7 oxidation states

Anomalous behaviour of fluorine
1. Small size
2. ionisation enthalpy, electronegativity, and
electrode potentials are higher.
3. Non availability of d orbitals in valence shell.
4. ionic and covalent radii, m.p. and b.p. ,low F-F
bond dissociation enthalpy and electron gain
enthalpy are quite lower than expected.

Reactivity towards hydrogen
1. Affinity for hydrogen decreases from fluorine to
iodine.
2. Acidic strength : HF < HCl < HBr < HI.
due to decrease in bond (H–X) dissociation
enthalpy
3. Reducing character :HF < HCl < HBr < HI.
due to decrease in bond (H–X) dissociation
enthalpy
4. Stability :due to decrease in bond (H–X)
dissociation enthalpy. H–F > H–Cl > H–Br > H–I.

Reactivity towards oxygen
1. Fluorine : OF2 and O2F2.
2. Chlorine, bromine and iodine form oxides in +1 to
+7
3. Chlorine oxides, Cl2O, ClO2, Cl2O6 and Cl2O7 .
ClO2used as bleaching agent for paper pulp and
textiles and water treatment.
4. Bromine oxides, Br2O, BrO2 , BrO3 .
5. Iodine oxides, I2O4 , I2O5, I2O7 .
I2O5 is very good oxidising agent and used in
estimation of carbon monoxide.27-Oct-19 Vijaykumar Nazare 87

Reactivity towards metals
• Metal halides:
Mg ( s ) + Br2 ( l ) → MgBr2 ( s )
• The ionic character decreases:MF >MCl >MBr >MI
(M is a monovalent metal)+1
• Halides in higher oxidation state – covalent .
• Eg. SnCl4, PbCl4, SbCl5 and UF6 are more covalent
than SnCl2, PbCl2, SbCl3 and UF4 respt.

1. Fluorine exhibits only –1 oxidation state
whereas other halogens exhibit + 1, + 3, + 5 and
+ 7 oxidation states also. Explain.
Solution
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 also.

Chlorine
• Chlorine (Greek, chloros = greenish yellow).
1. Preparation (oxidation by oxidising agent)
(i) By heating manganese dioxide with conc. hydrochloric acid.
MnO2 + 4HCl → MnCl2 + Cl2 + 2H2O
mixture of common salt and conc. H2SO4 is used in place of HCl.
MnO2 + 4NaCl + 4H2SO4 → MnCl2 + 4NaHSO4 + 2H2O + Cl2
(ii) By the action of HCl on potassium permanganate.
2KMnO4 + 16HCl → 2KCl + 2MnCl2 + 8H2O + 5Cl2
1. Manufacture of chlorine
(i) Deacon’s process: oxidation of hydrogen chloride gas by atmospheric
oxygen in the presence of CuCl2 (catalyst) at 723 K.
• (ii) Electrolytic process: electrolysis of brine (concentrated NaCl
solution). Chlorine is liberated at anode. It is also obtained as a by–
product in many chemical industries.27-Oct-19 Vijaykumar Nazare 90

Properties
1. greenish yellow gas , pungent and suffocating odour.
2. liquefied into greenish yellow liquid ,soluble in water.
3. Reaction with metals and non-metals form chlorides.
2Al + 3Cl2 → 2AlCl3 ; P4 + 6Cl2 → 4PCl3
2Na + Cl2 → 2NaCl; S8 + 4Cl2 → 4S2Cl2
2Fe + 3Cl2 → 2FeCl3 ;
4. affinity for hydrogen: reacts with compounds containing hydrogen to
form HCl.
H2 + Cl2 → 2HCl
H2S + Cl2 → 2HCl + S
C10H16 + 8Cl2 → 16HCl + 10C
5. Excess ammonia, chlorine gives nitrogen and ammonium chloride.
8NH3 + 3Cl2 → 6NH4Cl + N2
(excess)
• Excess chlorine : nitrogen trichloride (explosive) is formed
NH3 + 3Cl2 → NCl3 + 3HCl
(excess)

Properties
• Action on alkalies :
1. With cold and dilute alkalies : chloride and hypochlorite.
2NaOH + Cl2 → NaCl + NaOCl + H2O
(cold and dilute)
2. hot and concentrated alkalies :chloride and chlorate
6 NaOH + 3Cl2 → 5NaCl + NaClO3 + 3H2O
(hot and conc.)
3. Dry slaked lime : bleaching powder.
2Ca(OH)2 + 2Cl2 → Ca(OCl)2 + CaCl2 + 2H2O
4. composition of bleaching powder is
Ca(OCl)2.CaCl2.Ca(OH)2.2H2O.

Chlorination
Photochemical of hydrocarbon
Chlorine reacts with hydrocarbons and gives
substitution products with saturated hydrocarbons
and addition products with unsaturated
hydrocarbons.

Properties of Chlorine
1. Strong oxidising agent: 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 + H2O + Cl2 → Na2SO4 + 2HCl
SO2 + 2H2O + Cl2 → H2SO4 + 2HCl
I2 + 6H2O + 5Cl2 → 2HIO3 + 10HCl
2. powerful bleaching agent :bleaching action is due
to oxidation.
Cl2 + H2O → 2HCl + O
Coloured substance + O → Colourless substance

Uses
1. for bleaching woodpulp (required for the
manufacture of paper and rayon), bleaching cotton
and textiles .
2. extraction of gold and platinum
3. manufacture of dyes, drugs and organic compounds
such as CCl4, CHCl3, DDT, refrigerants, etc.
4. sterilising drinking water .
5. As disinfectant .
6. preparation of poisonous gases such as phosgene
(COCl2), tear gas (CCl3NO2), mustard gas
(ClCH2CH2SCH2CH2Cl).

Hydrogen Chloride
• Preparation :
In laboratory:Heating sodium chloride with
concentrated sulphuric acid.
• HCl gas is dried by passing concentrated sulphuric
acid.

Properties
• Colourless ,pungent smelling gas .
• Soluble : HCl(g) + H2O (l) → H3O + (aq) + Cl− (aq) Ka = 107
High (Ka) value - strong acid .
• Reaction with NH3 : white fumes of NH4Cl.
NH3 + HCl → NH4Cl
• aqua regia :three parts of conc HCl + one part of conc
HNO3
used : dissolving noble metals, e.g., gold, platinum.
Au + 4H+ + NO3
− + 4Cl− → AuCl−
4 + NO + 2H2O
3Pt + 16H+ + 4NO3 + 18Cl− → 3PtCl6
− + 4NO + 8H2O
• 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
1. Manufacture of chlorine, NH4Cl and glucose
(from corn starch)
2. extracting glue from bones and purifying
bone black
3. In medicine as laboratory reagent.

HALIC(I) ACID
(HYPOHALOUS
ACID)
HOF(HYPOFLUOR
OUS ACID)
HOCL(HYPOCHLOR
OUS ACID)
HOBR(HYPOBROMOU
S ACID)
HOI(HYPOIODOUS
ACID)
Halic (III)
acid(Halous acid) – HOCIO(chlorous acid) – –
Halic (V) acid(Halic
acid) – HOCIO2(chloric acid) HOBrO2(bromic acid) HOIO2(ionic acid)
Halic(VII)
acid(Perhalic acid) –
HOCIO3(perchloric
acid) HOBrO3(perbromic acid)
HOIO3(periodic
acid)
Oxoacids of halogens

Interhalogen Compounds
• Halogens react with each other due to electronegativity difference.
• More electronegative (smaller halogen)– anion
Less electronegative (higher halogen) - cation
• Types : XX’ , XX’3 , XX’5 and XX’7 where X - halogen of larger size ,
more electropositive and X’ – halogen of smaller size .
• Ratio of X and X’ increases ,number of atoms per molecule also
increases .eg. IF7
• Preparation : direct combination of halogen on lower interhalogen
compounds.

TYPE FORMULA PHYSICAL STATE AND COLOR STRUCTURE
XX’1
ClF
BrF
IFa
BrClb
ICl
IBr
colorless gas
pale brown gas
detected spectroscopically gas
ruby red solid(α-form)
brown red solid (β – form)
Black solid
-
-
-
-
-
-
XX’3
ClF3
BrF3
IF3
IClc
3
colorless gasyellow green liquidyellow powder
orange solid
Bent T-shaped
Bent T-shaped
Bent T-shaped(?)
Bent T-shaped(?)
XX’5
IF5
BrF5
ClF5
colorless gas but
solid below 77 K
colorless liquidsquare pyramidal
square pyramidal
square pyramidal
XX’7 IF7 colorless gas
pentagonal
bipyramidal

• Deduce the molecular shape of BrF3 on the basis of VSEPR theory.
Solution
The central atom Br has seven electrons in the valence shell. Three of
these will form electron- pair bonds with three fluorine atoms
leaving behind four electrons. Thus, there are three bond pairs and
two lone pairs. According to VSEPR theory, these will occupy the
corners of a trigonal bipyramid. The two 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. In addition, the axial fluorine atoms will be bent
towards the equatorial fluorine in order to minimise the lone-pair-
lone pair repulsions. The shape would be that of a slightly bent ‘T’.27-Oct-19 Vijaykumar Nazare 102

Properties of interhalogens
• Covalent and diamagnetic .
• Volatile solids or liquids except ClF gas .
• Inter halogen compounds are more reactive
due to week X-X’ bond than X-X bond (F-F)
• Hydrolysis :
XX’ + H2O → HX’ + HOX

Uses
• non aqueous solvents.
• fluorinating agents.
• ClF3 and BrF3 are used for the production of
UF6 in the enrichment of 235U.
U(s) + 3ClF3(l) → UF6(g) + 3ClF(g)

Group 18 Elements
• helium, neon, argon, krypton, xenon and radon.
chemically unreactive. They form very few
compounds - noble gases.
• Why are the elements of Group 18 known as
noble gases ?
Solution
valence shell orbitals completely filled .
react with a few elements only under certain
conditions.
Therefore, they are now known as noble gases.

1. Electronic Configuration : ns2np6 , helium 1s2
2. Ionisation Enthalpy : very high
Due to stable electronic configuration.
decreases down the group with increase in
atomic size.
3. Atomic Radii : increase down the group with
increase in atomic number.
Larger than group 17 due to e-e repulsion .
4. Electron Gain Enthalpy : large positive values .
stable electronic configurations, no tendency to
accept the electron .

Physical Properties
• Monoatomic , colourless, odourless and tasteless.
• sparingly soluble in water.
• low melting and boiling points : because
interatomic interaction in elements is weak
dispersion forces.
• Helium can diffuse through rubber , glass or plastic
.

Chemical Properties
1. Least reactive (inertness to chemical reactivity )
Due to
(i) The noble gases except helium (1s2 ) have completely filled
ns2np6 electronic configuration in their valence shell.
(ii) high ionisation enthalpy and more positive electron gain
enthalpy.
2. (a) Xenon-fluorine compounds
Xenon forms three binary fluorides, XeF2, XeF4 and XeF6
XeF6 can also be prepared by the interaction of XeF4 and O2F2 at
143K.
XeF4 + O2 F2 → XeF6 + O2
XeF2, XeF4 and XeF6 are colourless crystalline solids.
3. powerful fluorinating agents.
4. Hydrolyses: XeF2 is hydrolysed to give Xe, HF and O2.
2XeF2 (s) + 2H2O(l) → 2Xe (g) + 4 HF(aq) + O2(g)

Structures

Uses
1. Helium: non-inflammable and light gas. filling balloons for
meteorological observations.
2. used in gas-cooled nuclear reactors.
3. Liquid He : cryogenic agent .
4. produce and sustain powerful superconducting magnets for
NMR and MRI .
5. diluent for oxygen in modern diving apparatus because of its
very low solubility in blood.(scuba divers)
6. Neon: discharge tubes and fluorescent bulbs for advertisement
display purposes. Neon bulbs :botanical gardens and green
houses.
7. Argon: inert atmosphere for metallurgical processes (arc
welding of metals or alloys) and for filling electric bulbs.
8. laboratory :handling substances that are air-sensitive.
9. Xenon and Krypton : light bulbs designed for special purposes.

P block elements

Recommended

Recommended

More Related Content

What's hot

What's hot (20)

Similar to P block elements

Similar to P block elements (20)

Recently uploaded

Recently uploaded (20)

P block elements