7. ELEMENTS OF GROUPS 16, 17 AND 18
 Se, Ge, As, Sb, Te are metalloids
 Al, Ga, In, Tl, Sn, Pb, Bi, Po are metals
 Po, At, Rn are radioactive elements
 Metal have tendency of donating, non metal have tendency of
accepting and metalloids have both the property
 Last electron in p subshell
 Grp 13 to Group 18 are p block
Electronic configuration of elements of group 16, 17 and 18:
OXIDATION STATE :
Group
13
14
15
16
17
18
Electronic
configuratio
n
ns2 np1
ns2 np2
ns2 np3
ns2 np4
ns2 np5
ns2 np6
Valence
electron
Group
oxidation
state or
higher
oxidation
state
Lower
oxidation
state
INERT PAIR EFFECT :
Due to poor screening or poor shielding of D and F orbital, Z effective
incresses and force of attraction incresses because of Z effective
nucleus hold ns2 pair tightly and make it inert
ns2 np6
DOWN THE GROUP, ns2 PAIR BECOMES INERT
 Stability of lower oxidation state :
 Stability of lower oxidation state :
ANOMALOUS BEHAVIOUR OF FIRST MEMBRANE
 Small size
 High electronegativity
 High ionisation energy
 Absence of d-block bcos 1st membrane are from 2nd period so there is no d
subshell
Atomic and physical properties of elements of group 13,
14, 15, 16, 17 and 18.
Atomic properties of Group 13, 14, 15, 16, 17 and 18 elements :
 Atomic size/ Atomic radii
In group 16, 17 and 18 atomic and ionic radii increase down the group, bcos of
increase in the number of quantum shells
Across a period atomic or ionic radii decrease with increasing atomic number,
consequent to increase in (Zeff ) effective nuclear charge. Group 17 elements
(Halogens) have the smallest atomic radii in their respective periods
ii. IONISATION ENTHALPY :
 The ionisation enthalpy decreases down the group due to increase in the
atomic size.
 Across a period ionisation enthalpy increases with increase of atomic number.
This is due to addition of electrons in the same shell. However the elements of
group 16 have lower ionisation enthalpy compare to group 15 bcos of half filled
electronic configuration so it become more stable than group 16 bcos of this it
required more IE to remove electron from grp 15
ELECTRON AFFINITY/ ELECTRON GAIN ENTHALPY :
• In the groups 16 and 17 electron gain enthalpy becomes less negative
down the group.
• however in group 16, oxygen has less negative electron gain enthalpy
than sulfur due to its small atomic size.
• In group 17, fluorine has less negative electron gain enthalpy than
that of chlorine. This is due to small size of fluorine atom.
• Group 18 elements (noble gases) have no tendency to accept
electrons because of their stable electronic configuration (ns2np6)
and thus have large positive electron gain enthalpy.
• Highest electron gain enthalpy = Cl
• lowest (most positive) electron gain enthalpy in the periodic table is
helium (He) — and in fact, all noble gases (Group 18) have highly
positive electron gain enthalpies.
 If we give electron to N atom we have to give more energy , so this process
known as endothermic deltaHeg = +ve
 But if we give electron to Cl atom it required less energy , so this process is
known as exothermic deltaHeg = -ve
iii. ELECTRONEGATIVITY :
• In a group (16, 17 and 18) the electronegativity decreases down
the group
• F atom have highest EN and than N atom
• Halogens have very high electronegativity. Fluorine is the most
electronegative element in the periodic table.
GROUP 16 / OXYGEN FAMILY / CHALCOGEN FAMILY are ore
forming elements ( oxides or sulphide )
non-metals
Metalloids
Metal and radioactive
• Oxygen is the most abundant of all the elements on earth. Oxygen forms
20.95 % by volume of air and about 46.6 % by mass
• of earth's crust.
• Sulfur forms 0.034% by mass of the earths crust
• Sulphur mainly exists as sulphides and sulphates such as
gypsum (CaSO4 .2H2 O),
epsom salt (MgSO4 .7H2 O),
baryte (BaSO4 )
and sulfides such as galena (PbS),
zinc blende (ZnS),
copper pyrites (CuFeS2 ).
• Selenium and tellurium are also found as metal selenides and tellurides in
sulfide ores.
• Polonium which is radioactive is a decay product of thorium and uranium.
MELTING POINT AND BOILING POINT
 MP and BP are directly proportional to the mass
 MP and BP incresses from oxygen to polonium due to incresses
in the magnitude of vander waal’s force of attraction
M.P
O < S < Se < Po < Te
B.P
O < S < Se < Po < Te
Oxidation State ns2np4
Tendency to form -2 O.S directly proportional to EN
O > S > Se > Te > Po
Tendency to form +4O.S
O < S < Se < Te < Po
Tendency to form +6O.S
O > S > Se > Te > Po
Anomalous behaviour of oxygen :
i. Atomicity : Oxygen is a diatomic molecule (O2 ) while others are
polyatomic molecules. For example P4 , S8
ii. Magnetic property : Oxygen is paramagnetic (unpaired electron)
while others are diamagnetic (paired electron).
iii. Oxidation state : Oxygen shows -2, -1, and +2 oxidation states
while other elements show, -2, +2, +4, +6 oxidation states. Oxygen
can not exhibit higher oxidation state due to absence of vacant d
orbitals.
iv. Nature of hydrides : Hydride of oxygen (H2 O) is liquid at room
temperature bcos of hydrogen bond while hydrides of other
members of the group are gases. Bcos of weaker or absent
hydrogen bond
v. Common covalency of oxygen is 2. In rare cases it is four. But for the
other members of the group 16 the covalency can exceed four. The
anomalous behaviour of oxygen is due to the following reasons. i.
small atomic size ii. high electronegativity. iii. absence of inner dorbitals.
Chemical Reactivity towards hydrogen: i. Group 16 elements :
• The elements of group 16 react with hydrogen to form hydrides of
the type H2 E. (Where E = O, S, Se, Te, Po).
• H2 O, H2 S, H2 Se, H2 Te and H2 Po.
• H2 O is a colourless, odourless liquid, while H2 S, H2 Se, H2 Te and
H2 Po are colourless bad smelling, poisonous gases at ambient
conditions.
 Bond dissociation Enthalpy: H and O bond strong bcos
of Hydroden bond and O atom is small so more BDE but
in H and Po have weak bond so less BDE
H2O > H2S > H2Se > H2Te > H2Po
 Thermal stability: directly proportional to BDE
H2O > H2S > H2Se > H2Te > H2Po
 Acidic Character: depend on H+ removal
H2O < H2S < H2Se < H2Te < H2Po
 Reducing Character: depend on H
H2O < H2S < H2Se < H2Te < H2Po
 Bond length :
H2O < H2S < H2Se < H2Te < H2Po
 VOLATILE NATURE : H₂O is least volatile despite being light.
Because it has strong hydrogen bonding, holding the molecules
tightly together
H2S > H2Se > H2Te > H2Po > H2O
 COVALENT CHARACTER : Covalent character increases down the
group due to decreasing electronegativity and larger atomic size,
which results in weaker polarity of the H–E bond.
H2O < H2S < H2Se < H2Te < H2Po
 BOND ANGLE : As we go down the group the central atom become
larger and lone pair and bond pair repln decresses in result bond
angle get smaller
H2O > H2S > H2Se > H2Te > H2Po
Chemical Reactivity towards hydrogen:
H + E ------------> H2E (Where E = O, S, Se, Te, Po).
(H2O) hydrides of
oxygen
colourless, odourless
liquid,
Other element of group
16
H2S, H2Se, H2Te, H2Po
colourless bad smelling,
poisonous gases at
ambient conditions.
• All hydrides have angular structures which involve sp3 hybridization
Reactivity towards oxygen :
E + O -----> EO2
E + O -----> EO3
Ozone (O3 )
EO2
sulfur dioxide (SO2) selenium dioxide (SeO2 )
are gases,
solid.
SO2 is reducing while TeO2 serves as an oxidising agent.
EO3
SO3 , SeO3 , TeO3 are also acidic in nature.
Reactivity towards halogens :
E + X -----> EX2
E + X -----> EX4
E + X -----> EX6 (Where E = S, Se, Te)
Covalent character directly proportional to charge
EX2 < EX4 < EX6
STABILITY OF HALIDES
F > Cl > Br > I
Hexahalide
Sf6, Sef6, Tef6
Colourless
Sp3d2 hybridization
Octahedral geometry
Sf6 is more stable
than any other
hexaflouride
Tetrahalide
Sf4, Sef4, Tef4, TeCl4
Sp3d hybridization
Trigonal bipyramidal
geometry with one
equatorial position
occupied by a lone
pair.
Dihalides,
SCl2 , SeCl2 , TeCl2
sp3 hybridisation
tetrahedral structure
with two equatorial
positions occupied
by lone pairs.
Monohalides are dimeric in nature. For example, S2 F2 , S2 Cl2 ,
Se2 Cl2 and SeBr2 .
dimeric halides undergo disproportionation.
Reactivity towards metal
4Al + 3O2 -----------> 2Al2O3
4Li + O2 -------------> 2Li2O ( oxcite )
2Na + O2 ------------> Na2O2 ( peroxide )
4K + 3O2 -----------> KO2 ( superoxide )
2CU + S -----------> Cu2S
2Na + S ------------> Na2S
Allotropy :
At high temp
(1000k) S2 is
dominant
• Oxoacids of sulfur : Sulfur forms a number of oxoacids. Some of them
are unstable and cannot be isolated. They are known to exist in
aqueous solutions or in the form of their salts
• Some important oxoacids of sulfur and their structures are given
below.
Dioxygen Preparation
Decomposition of oxygen rich compound such as
i. Laboratory methods :
2KNO3(s) --------------> 2KNO2 + O2(g)
2H2O2 -----------------> 2H2O + O2(g)
ii. Electrolysis :
hydrogen is liberated at cathode and oxygen at anode.
iii. Industrial method : Dioxygen is obtained from air, by first removing
carbon dioxide and water vapour. The remaining gases are liquified
subsequently. This is followed by fractional distillation which gives
dinitrogen and dioxygen.
Physical properties :
• Dioxygen is colourless and odourless gas.
• Dioxygen is sparingly soluble in water, 30.8 cm3 of O2
dissolves in 1000 cm3 of water at 293 K. A small amount of
dissolved dioxygen is sufficient to sustain marine and aquatic
life.
• It liquifies at 90 K and freezes at 55 K.
• Oxygen has three stable isotopes 16O, 17O and 18O.
• Molecular oxygen, O2 paramagnetism.
Po short lived (half life 13.8 days)
Group 16 exhibit allotropy.
Chemcial Properties :
i. Reaction with metals : Dioxygen directly reacts with almost all
metals except Au, Pt to form their oxides.
4Li + O2 ------> 2Li2O
4Fe + O2 -------> 2Fe2O3
ii. Reaction with nonmetals :
Dioxygen reacts with nonmetals (except noble
gases) to form their oxides.
C + O2 -------> CO2 ( O2 in excess )
C + O2 -------> CO
S + O2 -------> SO2
Oxides can be classified into
a. Acidic oxides
b. Basic oxides
c. Amphoteric oxides
d. Neutral oxides
a. Acidic oxides : An oxide which react with water to give an acid or
reacts with a base to give a salt is called acidic oxide.
For example, SO2 , SO3 , CO2 , N2 O5 , Cl2 O7 etc.
Generally non metals
CO2 + H2O --------> H2CO3 (CARBONIC ACID)
N2O5 + H20 ---------> 2HNO3 (NITRIC ACIID)
Cl2O7 + H2O ------------> 2HClO4 ( very strong acid )
b. Basic oxides : An oxide which dissolves in water to give a base
or reacts with an acid to give salt is called basic oxide.
For example, Na2 O, CaO, BaO etc.
Generally metals
Na2O + H2O ------> 2NaOH
K2O + H2O --------> 2KOH
CaO + 2HCl ---------> CaCl2 + H2O
c. Amphoteric oxides : The oxide which reacts with a base as
well as with an acid to give salt is called an amphoteric oxide.
For example, Al2 O3
d. Neutral oxides : The oxides which are neither
acidic nor basic, are called as neutral oxides.
For example, CO, NO, N2 O etc.
Ozone : Ozone (O3 ) is an allotrope of oxygen. Oxygen in
the upper atmosphere absorbs energy in the form of
ultra-violet light and changes to atomic oxygen, which
combines with molecular oxygen to form O3 .
a. Preparation of Ozone :
• Ozone is prepared in the laboratory by passing silent electric
discharge through pure and dry oxygen in an apparatus called
ozoniser. As the conversion of oxygen to ozone is only 10%,
the product is known as ozonised oxygen. It is an
endothermic process
b. Physical properties of ozone :
i. Pure Ozone is a pale - blue gas, dark blue
liquid and violet black solid.
ii. Ozone has a characteristic smell. When
inhaled in concentration above 100 ppm, it
causes nausea and headache.
iii. It is diamagnetic in nature.
Chemical Properties :
i. Oxidising property : Ozone is a powerful oxidising agents
as it easily decomposes to liberate nascent oxygen.
(O3 ----------> O2 + O).
This rxn is responsible for depletion of
ozone layer bcos NO is present in exhaust
in jet plane or normal plan
What is Ozone Depletion?
•Thinning of the ozone (O₃) layer in the upper atmosphere (stratosphere).
☀️ Role of Ozone
•Ozone absorbs harmful UV radiation from the sun, protecting life on Earth.
⚗️ Causes
•Nitrogen oxide (NO) from car exhausts and jet aircrafts reacts with ozone:
•Chlorofluorocarbons (CFCs) from aerosols and refrigerators also destroy
ozone.
📍 Affected Areas
•Most severe depletion occurs over polar regions, especially Antarctica.
🚨 Impact
•More UV rays reach Earth, leading to: incresesd skin cancer, eye cataracts and
genetic as well as immune system damage among people
d. Structure of Ozone : Ozone (O3 ) is an angular molecule. The two
O---O bond lengths in the ozone molecule are identical, 128 pm and
the O---O---O bond angle of about 117°. It is a resonance hybrid of two
canonical forms.
Platinsed asbestosis
Sulfuric acid, H2 SO4
a. Preparation : Sulfuric acid is manaufactured by
Contact process, which involves the following three
steps
i. Sulfur or sulfide ore (iron pyrites) on burning or
roasting in air produces sulfur dioxide.
• The reaction is exothermic and reversible
• low temperature (720K) and high pressure (2 bar) are
favourable conditions for maximum yield of SO3 .
b. Physical properties of H2SO4 :
i,. Sulfuric acid is a colourless, dense, oily liquid.
ii. It has a density (specific gravity) of 1.84 g/cm3 at
298 K
iii. It freezes at 283 K and boils at 611 K.
iv. It is highly corrosive and produces severe burns on
the skin.
v. Strong oxidizing agent and dehydrating agent
d. Uses : Sulfuric acid is a very important industrial chemical. It
is used
• In the manufacture of fertilizers. For example, ammonium
sulfate, superphosphate, etc.
• In the manufacture of pigments, paints and dyestuff
intermediates.
• In petroleum refining.
• In detergent industry.
• In metallurgy, for cleaning of metals electroplating and
galvanising.
• In storage batteries. • As a laboratory reagent.
• In the manufacture of nitrocellulose products
GROUP 17/HALOGENS FAMILY (nS24p5)
(Greek halo means salt, gene means born), that is, salt
producing element.
F
Cl
GAS
LIQUID
Br
SOLID
I
RADIOACTIVE
Halogens are very reactive due to high electronegativities and hence
they are not found in free sate. They occur in the form of compounds.
🌟 Natural Occurrence of Halogens
🟦 Fluorine (F)
•Mostly found as insoluble fluorides:
• CaF₂ (fluorspar)
• Na₃AlF₆ (cryolite)
• Fluorapatite: 3Ca3(PO4)2⋅CaF2
• Also present in:
• Soil and freshwater plants
• Animal bones and teeth (essential for dental health!)
🌊 Chlorine (Cl), Bromine (Br), and Iodine (I)
•Found in seawater mostly as salts:
• NaCl (dominant: 2.5% by mass)
• Other chlorides, bromides, and iodides of Na, K, Mg, Ca
•Dried sea beds yield:
• Sodium chloride
• Carnallite: KCl⋅MgCl2⋅6H2O
🌱 Iodine (I)
•Concentrated in marine organisms:
• Seaweed: ~0.5% iodine
• Chile saltpetre: ~0.2% sodium iodate
☢️ Astatine (At)
•Rare and radioactive
•Half-life: 8.1 hours
•Due to its short half-life, it's only found in trace
amounts from radioactive decay
Physical properties of group 17 elements
Fluorine, chlorine are gases,
bromine is a liquid
iodine is a solid at room temperature.
F2 is yellow,
Cl2 greenish yellow,
Br2 red and
I2 is violet, in colour
Fluorine and chlorine react with water.
Bromine and iodine are only sparingly soluble in water and are soluble in various
organic solvents such as
chloroform,
carbon disulfide,
carbon tetrachloride,
hydrocarbons which give coloured solutions.
Bond dissociation enthalpies
Anomalous behaviour of fluorine :
i. small atomic size
ii. high electronegativity
iii. absence of d-orbitals in valence shell
iv. low F-F bond dissociation enthalpy
Some anomalous properties of fluorine :
i. Ionisation enthalpy, electronegativity, electrode potential are
all higher for fluorine than expected trends shown by other
halogen
ii. Ionic and covalent radii, m.p., b.p. and electron gain enthalpy
are quite lower than expected.
iii. Most of the reactions of fluorine are exothermic (due to the
short and strong bond formed by it with other elements).
iv. It forms only one oxoacid (HOF) while other halogens form a
number of oxoacids.
v. Hydrogen fluoride is a liquid (b.p. 293K) due to strong
hydrogen bonding while other hydrogen halides are gases.
Only HF dimerises to give H2F2
Chemical Properties of elements of groups 17
• All halogens exhibit -1 oxidation state. However Cl, Br and I exhibit
+1, +3, +5 and +7 oxidation states as well.
• This is because they are less electronegative than F and possess
empty d-orbitals in the valence shell and therefore, can expand the
octet.
• The oxidation states +4 and +6 occur in the oxides and oxoacids of Cl
and Br
• fluorine being most electronegative exhibits mostly -1 oxidation state
Chemical Reactivity towards hydrogen:
Acidic strength
It is due to decreasing bond dissociation enthalpy
Bond strength
Reducing character
Reactivity towards oxygen (Group 17)
O2F2 = unstable
OF2 = relatively stable at 298K
Chlorine oxides, Cl2 O, ClO2 , Cl2 O6 and Cl2 O7
are highly reactive oxidising agents and tend to explode.
O2 F2 oxidises plutonium to PuF6 and the reaction is used in
removing plutonium as PuF6 from spent nuclear fuel.
• ClO2 is used as bleaching agent for paper pulp and textiles and in
water treatment.
• Bromine oxides, Br2 O, BrO2 , BrO3 (are the least stable halogen
oxides (middle row anomally). They are very powerful oxidising
agents)
• Iodine oxides, I2 O4 , I2 O5 and I2 O7 are insoluble solids and
decompose on heating.
• I2O5 is a very good oxidising agent and used for the estimation of
carbon monoxide.
• Strong oxidizing agent
Reactivity towards halogens :
XX', XX'3, XX'5, XX’7,
Where X is the halogen atom with larger size and X' , is the halogen
atom with smaller size.
Reactivity towards metals : Group 17 elements
react with metals instantly to give metal halides.
• The metal halides having metals in their higher oxidation states are
more covalent than the ones having metals in lower oxidation state.
• For example, SnCl4 , PbCl4 , SbCl5 and UF6 are more covalent than
SnCl2 , PbCl2 , SbCl3 and UF4
Oxoacids of halogens :
• Only four oxoacids have been isolated in pure form:
hypofluorous acid (HOF),
perchloric acid (HClO4 ),
iodic acid (HIO3 ),
metaperiodic acid (H2IO6 ).
The others are stable only in aqueous solutions or in the form of their
salts
• The acid strength of the halogen oxoacids increases with the
increasing oxidation state of halogen.
• For example, acid strength increases from HClO, a weak acid
(Ka = 3.5 × 10-8), to HClO4 , a very strong acid (Ka >>1).
Chlorine and compounds of chlorine
• Chlorine was discovered by Scheele, a German Swedish chemist in
1774 by the action of HCl on MnO2 .
• In 1810, Davy established its elementary nature and suggested the
name chlorine on account of its colour.
(Greek, Chloros = greenish yellow).
Preparation :
2.
b. Manufacture of chlorine :
i. Deacon’s process : Chlorine is manufactured by the oxidation of
hydrogen chloride gas by atmospheric oxygen in the presence of
CuCl2 as catalyst at 723 K.
ii. Electrolytic process : By the electrolysis of brine (concentrated NaCl
solution), chlorine is liberated at the anode.
Uses : Chlorine is used
• For purification (sterilizing) of drinking water.
• For bleaching wood pulp required for manufacture of paper and
rayon, bleaching cotton and textiles.
• For extraction of metals like gold and platinum.
• In the manufacture of dyes, drugs and organic compounds such as
CCl4 , CHCl3 , DDT, refrigerants, etc.
• In the preparation of poisonous gases such as phosgene (COCl2 ),
tear gas (CCl3 NO2 ), mustard gas (ClCH2 CH2 SCH2 CH2 Cl).
Hydrogen Chloride : Hydrogen chloride was prepared by Glauber in
1648 by heating common salt with concentrated sulfuric acid. Davy in
1810 showed that it is a compound of hydrogen and chlorine.
Interhalogen compounds
• We have seen that all halogen molecules are diatomic.
• They form binary compounds with hydrogen, with oxygen they form
oxyacids, and with halogens they form interhalogen compounds.
• Due to this difference in electronegativity two or more halogen
atoms combine to form species which may be ionic or neutral.
• The neutral molecules are called Interhalogen compounds. For
example, CIF, BrF3 .
X is the halogen having larger size
and is more electropositive. X' is
the halogen having smaller size and
more electronegativity.
General characteristics interhalogen compounds
1. of The compound is considered as the halide of X. For example,
ClF. Here the halogen having larger size is chlorine, it is more
electropositive than F and hence the interhalogen compound is
named as chlorine monofluoride. (n) is the number of atoms of X'
attached to X
As the ratio [radius of X : radius of X'] increases the value of n
also increases.
2. Interhalogen compounds have even number of atoms 2, 4, 6,
8. For example, ClF3 has 4 atoms.
3. The properties of interhalogen compounds are generally
intermediate between those of the halogens from which they
are made.
4. The central halogen exhibits different oxidation states in
different interhalogen compounds.
5. Number of X' atoms in the compounds is always odd.
6. They are all diamagnetic .
GROUP 18 (NOBLE OR INERT GAS) Ns2Np6
helium (2 He),
neon (10 Ne),
argon (18 Ar),
krypton (36 Kr),
xenon (54 Xe),
radon (86 Rn)
🟦 Natural Distribution of Noble Gases
🟦 Presence in the Atmosphere
•All noble gases except radon are found in the air we breathe.
•Collectively, they make up around 1% by volume of dry air.
•Argon (Ar) is the most abundant of these in the atmosphere.
🔥 Sources of Helium and Neon
•The main commercial source of helium (He) is natural gas, especially in
deposits that contain high concentrations.
•Helium and neon (Ne) are also found in radioactive minerals like:
• Pitchblende , Monazite , Cleveite
•These minerals release noble gases during radioactive decay processes.
💎 Rarest Noble Gases
•Xenon (Xe) and radon (Rn) are the least abundant in the group.
c. Group 18 elements (Noble gases) :
 Noble gases are monoatomic.
 They are sparingly soluble in water.
 Noble gases have very low melting and boiling
points.
 Helium has the lowest boiling point (4.2 K) of any
known substances.
Group 18 elements (noble gases)
• electronic configuration ns2np6
• zero valent and mostly exist as monoatomic gases.
• xenon exhibits higher oxidation states. Its outermost
shell has d-orbitals. The paired electrons of the valence
shell can be unpaired and promoted to empty d-orbitals.
• higher oxidation state such as XeF2 (+2), XeF4 (+4), XeF6
(+6), XeO3 (+6) and XeOF4 (+6) are formed.
Chemical Reactivity towards hydrogen: Group 18 elements (Noble
gases)
• Noble gases are chemically inert towards hydrogen due to their
stable electronic configuration.
Reactivity towards oxygen : Group 18 elements
• Noble gas elements are chemically inert and do not directly
react with oxygen
Reactivity towards halogens :
Reactivity towards metals : Group 18 elements
Noble gases do not directly react with metals.