P BLOCK ELEMENT

P BLOCK ELEMENT
Group IIIA or 13
Group IIIA or 13 of the long form of the periodic table consists of five element – boron, aluminum gallium,
indium and thallium.
General Characteristics
Atomic and ionic radius: The atomic radius increases from boron to thallium. Ionic radius (M3+) also
increases.
Density: Density increases from boron to thallium. However, boron and aluminium have comparatively low
values.
Melting and boiling points : The melting point decreases from B to Ga and then increases.
Ionisation energy : It decreases from B to Al, increases from Al to Ga, again decerase from Ga to In and
again increases from in In to Tl .
Oxidation states: The expected oxidation states are +3 and +1. Boron shows +3 oxidation state in all its
compounds. Other members show +3 and +1 oxidation states. The stability of +1 oxidation state increases
from aluminium to thallium and the stability of +3 oxidation state decrease from aluminium to thallium.
Nature of bonding : Boron is considerably smaller than other members and its ionisation potential is
maximum and hence it always shows covalency. Many simple compounds of the other elements such as
AlCl3 and GaCl3 are covalent when anhydrous. However, Al, Ga, In and TI all form metal ions in solution.
The change from covalent to ionic occurs because the ions are highly hydrated and the amount of hydration
energy evolved exceeds the ionisation energy. Thallium compounds are ionic.
Compounds
Oxides and hydroxides: All the members of this group form oxides of the type M2O3. On moving down the
group there is a gradual change from acidic to amphoteric and then to basic character of the oxides.
B2O3
Al2O3
Ga2O3 In2O3
TI2O3.
More acidic Amphoteric Basic Less basic
Basic
All these elements form hydroxides of the type M(OH)3. The basic nature of the hydroxides increases down
the group the change from acidic nature to basic H3BO3 acts as an acid. B(OH)3 in aqueous medium
coordinates a molecular of water to form a hydrated species.,
H
O  B(OH)3.B3+ ion pulls the σ electron
H
Charge of the coordinated O-atom towards itself . the coordinated oxygen , in turn,pulls the σ electron
charge of the O―H bond of the attached water molecule towards itself. This facilitates the removal of H+
ion from the O―H bond. The aqueous solution acts as a weak acid.
H
O  B(OH)3  [B(OH)4]- + H+
H
Hydrides: Boron forms a number of stable covalent hydrides with a general formulae BnHn+4 and BnHn+6.
these are called boranes. The representative compounds of the two series are B2O6 (Diborane) and
B4H10 (Tetraborane-10). Boranes are electron deficient compound.
Aluminium forms a polymeric hydride of the formula. (AlH3)n, commonly known as alane, Gallium
forms a dimeric hydride Ga2H6 (digallane) and indium forms a polymeric hydride, (InH3)n. thallium
does not form any hydride.
Boron, aluminium and gallium form complex anionic hydrides such as:
NaBH4 Sodium borohydride.LiAlH4 Lithium aluminium hydride LiGaH4 Lithium gallium hydride These
are powerful reducing agents.
Halides: All the members of this group form trihalides. The boron halides are covalent. The boron atom in
BX3 molecule acquires six electrons in its outermost shell. It can accept a pair of electrons from a suitable
donor to attain an octet configuration. Thus, BX3 act as Lewis acids.
BX3 + NH3  {H3N  BX3}
All boron trihalides except boron trifluoride are hydrolysed to boric acid.
BCl3 
3HCl 


BBr3   3H2 O H3BO3  3HBr
3HI 
BI3 


The trifluorides of Al, Ga, In and TI are ionic while the chlorides, bromides and iodides are largely covalent
when anhydrous. However, their covalent nature decreases on moving from Ga to Tl. Trihalides fume in air
and undergo hydrolysis.
AlCl3 + 3H2O  Al(OH)3 + 3HCl
Halides of boron
Boron combines with halogens and forms the halides of type BX3, (X = F, Cl, Br, 1)
(i)
The trihalides are electron deficient compounds. boron atom acquires six electrons on account of
three B ─X bonds. Thus, the boron atom in BX3 molecule can accept two more electrons.
H3N + BF3 → [H3N → BF3]
Donor
Acceptor
(Lewis base) (Lewis acid)
The relative Lewis acid character of boron trihalides is found to follow the following order,
BI3 > BBr3 > BCl3 > BF3.
But the expected order on the basis of electronegativity of halogens decreases from F to I) should be,
BF3 > BCl3 > BBr3 > BI3.
This anomaly is explained on the basis of the relative tendency of the halogen atom to back donate its
unutilized electrons to vacant p-orbital of boron atom. In BF3, boron has a vacant 2p-orbital and each
fluorine has fully filled unutilized 2p-orbitals Fluorine transfers two electrons to vacant 2p-orbital of boron,
thus forming p - p bond.
This bond reduces the electron deficiency of boron atom hence its Lewis acid character decreases. The
tendency to form back bonding is maximum in BF3 and decreases from BF3 to BI3. Thus, BCl3, BBr3 and
BI3 are stronger Lewis acids than BF3.
Thermite wilding of metals : Another application is the welding of metals especially the welding of steel. A
mixture of aluminium powder and Fe2O3 in the ratio of 1 : 3 (known as thermite ) is taken in a crucible lined
with magnesite and having a plug hole .
Alums :(M2SO4 M’2(SO4)3.24H2O)
Where M stands for monovalent basic radicals such as Na+, K +, Rb+ , Ag+, Tl +, NH+4 and M’ for trivalent
basic radicals suchas Al3+ , Cr3+, Fe3+, Mn3+ Co3+, etc .
Some examples of alums are:
Potash alum
K2SO4 .Al2( SO4).24H2O
( Commonly called alum )
Ammonium alum
(NH4)2 SO4.Al2(SO4)3. 24H2O
Sodium alum
Na2SO4. Al2 (SO4)3 .24H2O
2
Chorme alum
Ferric alum
K2 SO4. Cr2 (SO4)3. 24H2O
(NH4)SO4.Fe2(SO4)3. 24H2O
THE ELEMENTS OF GROUP IVA OR 14(Elements of Carbon Family)
Element At.No. Electronic Configuration
Carbon 6
2, 4
1s2, 2s22p2
Inert gas core
[He], 2s22p2
Silicon
14
German 32
ium
50
Tin
82
Lead
2, 8, 4
1s2, 2s22p6, 3s23p2
[Ne] 2s2 3p2
2, 8, 18, 4
1s2, 2s2 2p6, 3s2 3p6 3d10, 4s2 4p2
[Ar] 3d10, 4s2 4p2
2, 8, 18, 18, 4
5s25p2
1s2, 2s2 2p6, 3s2 3p6 3d10, 4s2 4p6 4d10, [Kr] 4d10,5s2 5p2
2, 8, 18, 32, 18, 4 1s2, 2s2 2p6, 3s2 3p6 3d10, 4s2 4p6 4d10 4f14,
5s2 5p65d10, 6s2 6p2
Non-metallic and Metallic character : The change from non-metallic to metallic character with the increase
of atomic number is best illustrate by this group. Carbon, the first element, is a non-metal, silicon, the second
element has most of the properties of a non-metal. Germanium is a metalloid while last two members Sn and
Pb are distinctly metals.
Atomic raddi and atomic volume : Atomic raddi and atomic volume increase gradually on moving down
the group. The size increase due to the effect of extra shell being added from member to member.
Ionisation potential : The ionization potentials decrease gradually from carbon to lead but not
systematically.
Melting and boiling points : Carbon has an extremely high melting point ( 3723°C). Silicon melts
appreciably lower than carbon but the values of silicon and germanium are still high. (Si = 1420°C; Ge =
945°C). Tin and lead are metallic and have much lower melting points (Sn = 232°C; Pb = 327°C) because
the M—M bonds are weaker.
Electronegativity : The electronegativity values do not decrease in a regular manner.
Valency : All these elements show a covalency of 4. When ns2 electrons of outermost shell do not participate
in bonding it is called inert pair and the effect is called inest pair effect. The last three elements have a
tendency to form M2+ ions as well as M4+ ions. Since the inert pair effect increases from Ge to Pb, the stability
of M4+ ions decreases and that of M2+ ions increases. Thus, the stability of these inos
Ge2+ < Sn2+ < Pb2+
Ge4+ > Sn4+ > Pb4+
The compounds of Ge2+ are unstable while compounds of Ge4+ are stable. The compounds of Sn2+ are less
stable than Sn4+. In the case of lead, Pb2+ compounds are more stable than Pb4+ compounds. The Pb4+
compounds, thus, act as oxidizing agents.
Catenation : The linking of identical atoms with each other to form long chains is called catenation. All the
elements of this group have the property of catenation. How ever, this property decreases from carbon to
lead. The decrease of this property is associated with M—M bond energy which decreases from carbon to
lead.
3
Maximum covalency : Maximum covalency of carbon is four because it has no d-orbitals or vacant orbitals
which can be used to accommodate more electrons. The remaining elements, however, have vacant dorbitals. These permit the formation of coordinate bonds with orther atoms or ions having lone pairs of
electrons. For example, SiF4 can combine with 2F- ions.
SiF4 + 2F-  [SiF6] 2Thus, the maximum covalency of silicon can be 6.
Hydrides : All the members of this group form covalent hydrides of the type MH4. Besides MH4, carbon
forms a large number of hydrides, saturated as well as unsaturated. The hydrides of Mh4 type are gaseous
and their thermal stability decreases and consequently the reducing nature increases from top to bottom.
However, among the hydrides, silicon hydrides are least stable to hydrolysis.
Hydrides
CH4
SiH4
GeH4
SnH4
PbH4
Decomposition temp. (°C) 800
450
285
150
0
The low stability of GeH4, SnH4 and PbH4 is due to weak M—H bond. It is due to large difference in the size
of M and hydrogen atom leading to poor overlapping and weak covalent bond.
Halides : The members of this group form tetrahalides of the type MX4 except PbBr4 and PbI4. The halides
are covalent and formed by sp3 hybridization. The thermal stability of halides of different elements with a
common halogen decreases with increasing atomic number. The thermal stability of tetrahalides of the same
element decreases with increase in molecular mass of the tetrahalide.
CX4 > SiX4 > GeX4 > SnX4 > PbX4
CF4 > CCl4 > CBr4 > CI4
[Non-existence of PbBr4 and PbI4 can be explained on the basis of strong oxidizing nature of Pb4+. The ions
Br- and I- are reducing agents, i.e., in presence of these ions, Pb4+ ions are reduced to Pb2+ ions.
Except carbon halides, other halides are readily hydrolysed by water. The trend towards hydrolysis, however,
decreases down the group. The hydrolysis is due to utilization of d-orbitals to which water molecules can get
attached. The tetrahalides of Si, Ge, Sn and Pb can form hexahalo complexes like [SiF6]2-, [GeF6]2-, [GeCl6]2. {SnCl6]2-. [PbCl6]2- with the corresponding halide ions. Thus, tetrahalides of Si, Ge, Sn and Pb act as strong
Lewis acids.
Oxides : All the elements of this group form oxides of the type MO2
CO2
SiO2 GeO2 SnO2 PbO2
The acidic nature decrerases with increase of atomic number. CO2 and SiO2 are acidic while GeO2, SnO2
and PbO2 are amphoteric
Carbon :
Allotropic forms of carbon :
Carbon exists in two allotropic forms :
(a) Crystalline and ( b) amorphous. The crystalline forms are diamond and graphite while the amorphous
forms are coal, charcoal, lamp black, etc.
Carbon
4
Crystalline
Diamond
Amorphous (micro crystalline )
Graphite
Coal
Charcoal
Lamp
black
(a) Diamond : (i) It is the purest form of carbon.
(ii) It is the hardest natural substance known
In diamond, each carbon atom is in sp3 hybridized state and linked to four other carbon atoms tetrahedrally
by covalent bonds.
(b) Graphite : It has a two dimensional sheet structure. Each carbon atom is in sp2 hybridized state and is
linked to three other carbon atoms in hexagonal planar structure. As there are four valence electrons in each
carbon atom, after forming three C—C bonds, each carbon atom is left with one spare electron in its porbital. This electron then overlaps with each other to form a -bond . Hence, the C—C distance in graphite
is shorter (1.42 Å) than that of diamond (1.54 Å). The -electrons are free to move throughout the entire
layers, graphite is good conductor of electricity. The adjacent layers are held by weak van der Waals’ forces
and the distance between two layers is sufficiently large (3.4 A).
Graphite is thermodynamically more stable than diamond and its free energy of formation is 1.9 kJ less than
diamond.
C
Graphite 1600

 Diamond

5000060000atm.
Gaseous Fuels : Some of the important gaseous fuels are discussed below.
Producer Gas
Producer gas is mainly a mixture of carbon monoxide and nitrogen. An ideal sample contains mainly about
35% CO and about 65% nitrogen with a small amount of CO2 and methane. It has a low calorific value. The
low calorific value (103 B.Th. U. per cubic ft.) is due to the presence of large proportion of nitrogen in it.
Water Gas
Water gas is mainly a mixture of CO and H2. In general, water gas consists of 40 volumes of CO, 50 volumes
of H2, 5 volumes of CO2 and 4-5 volumes of nitrogen. The calorific value of water gas is fairly high (310
B.Th.U. per cubic ft). The flame of water gas is short and hot, it is thus used for welding purposes.
Coal Gas :
Coal gas is a mixture of hydrogen, methane, carbon mohoxide, ethylene, acetylene, carbon dioxide, nitrogen
and oxygen. The calorific value of coal gas is 450-650 B.Th.U. per cubic ft.
Oil Gas : Oil has is obtained by cracking of kerosene oil. It is a mixture of lower hydrocarbons, mainly CH 4,
C2H4, C2H2, etc.
Natural Gas : Natural gas is found in regions rich in petroleum. The gas consists of chiefly methane. The
approximate composition of a sample of natural gas is given below :
CH4
C2H6
C3H8
C4H10
N2
85%
9%
3%
1%
2%
Liquefied Petroleum Gas (L.P.G.) : The main constituents of LPG are n-butane, isobutene, butane and
propane. The average calorific value of LPG is 55 kJg-1 or 29780 kcal/m3
5
Gobar Gas or Bio-gas : Its main constitutent is methane (60-70%). Other gases present are CO, H2, CO2,
H2S, N2 etc.
Glass : Glass is a amorphous super-cooled soild solution of silicates and borates. Its composition is variable
as it is not a true compound. An approximate formula for ordinary glass may be given as ,
R2O . MO . 6SiO2
Where R = Na or K and M = Ca, Ba, Zn or Pb., SiO2 may be replaced by Al2O3, P2O5.
NITROGEN FAMILY (V A GROUP OR 15th GROUP ELEMENT)
VA group or 15th group of the extended form of the periodic table consists of five elements nitrogen (N),
phosphorus (P) , arsenic (As), antimony (Sb) and bismuth (Bi).
Similarities and gradation in physical properties :
(a) Metallic and non-metallic character: Metallic nature increases as the atomic number increases. N
and P are purely non metals while Sb and Bi are metals. Arsenic behaves as a metalloid
N
P
AS
SB
Bi

 




Metalliod
Non

metals
Metals

(b)
Physical state: Nitrogen , the first element, is a gas while phosphorus, the second member, though a
solid, can pass readily into vapour state. And the remaining elements are solids.
(c)
Atomic radii : Atomic radii increase with the increase of atomic number.
Element
N
P
As
Sb
Bi
Atomic radii(Å)0.74 1.10 1.21 1.41 1.52
(d)
Ionisation potential: The ionization potential decreases regularly on descending the group as atomic
radii increases.
Element
N
P
As
Sb
Bi
Ionisation
336.0 253.9 231.0 199.1 184.9
potential (Kcal/mol)
(e) Electrogenativity: Electronegativity decreases gradually on descending the group from N to Bi
Element
N
P
As
Sb
Bi
Electronegativity 3.0 2.1 2.0 1.9 1.9
(f)
Density : It increases gradually on descending the group
Element
N
P
As
Sb
Bi
Density (g/mL)
0.809 1.823 5.73 6.62 9.78
(g)
Catenation: N,P and As exhibit the property of catenation but this property is much less than IVA
elements.
(i)
Atomicity: Nitrogen is diatomic gaseous molecule at ordinary temperature. Phosphorus, arsenic and
antimony all exist as discrete tetratomic tetrahedral molecules, viz. P4, As4 and Sb4 angle between X
―X―X is 60ο . the p-p bonding in phosphorus, arsenic , etc, is not possible due to large size of
these atoms.
(i)
Oxidation states: The elements of this group have five electrons in their outer shell. They exhibit a
maximum oxidation state of +5 towards oxygen by using all the five electrons of outer shell. The
tendency of the pair of ns electrons to remain inert (the inert pair effect) increases with increase of
atomic number. The stability of +3 oxidation state increases and that of +5 oxidation state decreases
6
on moving down from N to Bi. Nitrogen and phosphors generally exhibit –3 oxidation state due to high
electronegativity and small size. Nitrogen forms nitride ion (N3-) .
Oxides: All these elements form oxides of the type X2O3,X2O4 and X2O5.
Element Nitrogen Phosphorus Arsenic Antimony Bismuth
Type of
Acidic nature
oxides
increases
X2O3
N2O3
P2O3
As2O3 Sb2O3
Bi2O3
X2O4
N2O4
P2O4
As2O4 Sb2O4
Bi2O4
X2O5
N2O5
P2O5
As2O5 Sb2O5
Bi2O5
Acidic nature decrease
Greater is the electronegativity more is the acidic character of its oxide. Stability of oxides of higher
oxidation states decreases with increasing atomic number. Thermal stability decreases in each series from N
to Bi-N2O3 is most stable oxide.
(b) Oxyacids: All the elements of this group form oxyacids. Nitrogen forms a number of oxyacids but two
common oxyacids are nitrous acid (HNO2) and nitric acid (HNO3). Nitric acid is a stable acid while
HNO2 is unstable. Phosphorus forms a large number of oxyacids. The important ones are:
H3PO2
H3PO3
H4P2O4
Hypophosphorus Phosphorus
Hypophosphoric
Acid
acid
acid
H3PO4
HPO3
H4P2O7
Orthophosphoric Metaphosphoric Pyrophosphoric
Acid
acid
acid
Arsenic forms two oxyacids, H3AsO3 (arsenius acid) and H3AsO4 which exists in solution.
The strength and stability of oxyacids having the element in the same oxidation state decreases
gradually with decrease in electronegativity of central atom.
HNO3
H3PO4
H3AsO4
H3SbO4
Strong Weak
Very weak
Weakest
(c)
Hydrides: All the elements of this group form hydrides of the type MH3.
NH3
PH3
AsH3
SbH3
BiH3
Ammonia
Phosphine Arsine
Stibene
Bismuthine
(i)
All the hydrides are formed by the action of water or dilute acids on binary metallic compounds such
as Mg3N2, Ca3P2,Zn3As2,Mg3Sb2 and Mg3Bi2.
(ii)
The poisonous nature increases from NH3 to BiH3.
(iii) NH3 is highly soluble in water but other hydrides are less soluble.
(iv) The basic character decreases from NH3 to BiH3. NH3 is the strongest electron pair donor due to its
small size as the electron density of the electron pair is concentrated over a small region.
(v)
Thermal stability decreases gradually from NH3 to BiH3. the decomposition temperatures decrease
from NH3 to BiH3.
NH3
PH3
AsH3
SbH3
BiH3
1300°C
440 °C
280°C
150°C
Room temp.
(vi) The reducing nature increases. This shows that bond strength M―H decreases as electronegativty of
M decreases. NH3 is a weak reducing agent while AsH3, SbH3 and BiH3 are powerful reducing agents.
(vii) The shape of these hydrides is pyramidal. The formation is due to sp3 hybridization of central atom.
7
NH3
PH3
AsH3
SbH3
106.5°
93.5°
91.5°
91.3°.
(viii) hydrogen bonding : Hydrogen bonding is present in NH3 as the electronegativity difference between
nitrogen and hydrogen is high and the N – H bond shows polarity.
(d)
Halides: The main halides formed by the elements of this group are listed below:
Element
Trihalides
Pentahalides
Nitrogen
NF3,NCl3,NBr3,NI3
Phosphorus
PF3,PCl3,PBr3,PI3
PCl5,PF5,PBr5
Arsenic
AsF3,AsCl3,AsBr3,AsI3
AsF5
Antimony
SbF3,SbCl3,SbBr3,SbI3
SbF5,SbCl5
Bismuth
BiF3,BiCl3,BiBr3,BiI3
BiF5.
(i)
Trihalides : The elements directly combine with halogens and form trihalides, MX3. All the trihalides
are stable except NCl3, NBr3 and NI3. The unstable nature of NCl3, NBr3 and NI3 is due to low polarity
of N – X bond and a large difference in the size of nitrogen and halogen atoms. Due to the presence
of a lone pair of electrons on the central atom, they act as lewis bases. In the case of nitrogen halides,
the tendency to act as lewis bases decreases from NI3 to NF3
NI3 > NHr3 > NCl3 > NF3.
This is due to increased electronegativity from I to F. Also the tendency to act as Lewis base
decreases from N to Bi for a given halide.
The trihalides of phosphorus and antimony especially fluorides and chlorides act as Lewis acid also
by using the vacant d-orbitals.
Except NF3 and PF3, all the trihalides are hydrolysed by water.
NCl3+3H2O  NH3+3HClO
PCl3+3H2O  H3PO3+3HCl AsCl3+3H2O  H3AsO3+3HCl
SbCl3+H2OSbOCl+2HCl
BiCl3+H2O  BiOCl+2HCl
(ii)
Pentahalides: Bismuth does not form pentahalides because of inert pair effect.
PCl5  PCl3 + Cl2
PCl5 + H2O  POCl3 + 2HCl
POCl3 + 3H2O  H3PO4 + 3HCl
[PCl5 acts as an effective chlorinating agent. ]
X-ray studies have shown that the solid PCl5 is an ionic compound composed of [PCl4]+ [PCl6]-. Solid
PBr5 exists as [PBr4]+ Br-.
Nitric Acid HNO3
(a) Ostwald’s Process (Modern Proces)
Principle: The mixture of ammonia and air when passed over platinum gauze catalyst at 750-9000C,
the ammonia is oxidised to nitric oxide
Pt

4NH3 + 5O2 750
4NO+ 6H2O +21,600 calories
900
Chemical properties
Summary :
Concentration of nitric acid
Very Dilute NHO3
Dilute HNO3
Metal
Mg,Mg
Fe,Zn,Sn
Pb,Cu,Ag,Hg
Fe,Zn,
8
Main product
H2 + metal nitrate
NH4NO3 + metal nitrate
NO + metal nitrate
N2O + metal nitrate
Sn
Zn,Fe,Pb,Cu,Ag
An
NH4NO3 + Sn(NO3)2
Conc. HNO3
NO2 + metal nitrate
NO2 + H2SnSO3
Metastannic acid
1. Metals which do not react: Noble metals like good, platinum, iridium, rhodium, etc, are not acted
upon by nitric acid. However, these metals dissolve in aqua regia forms nascent chlorine which
attacks these metals.
THE ELEMENTS OF GROUP VIA OR 16 (Elements of Oxygen Family)
Group 16 or VIA of the extended form of periodic table consists of five elements - oxygen (O), sulphur (S),
selenium (Se), tellurium (Te) and polonium (Po). These (except polonium) are the ore forming elements and
thus called chalcogens,
2.
Physical characteristics :
(a) Physical state: Oxygen is a gas while others are solids. Oxygen molecule is diatomic while the
molecules of other elements are more complex. Sulphur, selenium and tellurium exist as staggered 8atom rings. However, the tendency to exist in 8-atom rings is maximum with sulphur and decreases
as we go down the group.
(b)
Metallic and non-metallic character: Metallic charcter increases with the increase of atomic
number.
(c)
Atomic radius, atomic volume and density: The values of these properties increase gradually as
expected with increase in atomic number.
O
S
Se
Te
Atomic radii.
0.74
1.04
1.14
1.37
Atomic volume
14.0
15.5
16.5
20.5
Density in solid state
1.14
2.07
4
6.23
(d)
Ionisation potentials: The ionization potentials are high and thus the elements do not lose the
electrons to form positive ions easily. The values decrease as the atomic number increases from O to
Po and thus the tendency to form positive ion increases gradually.
O
S
Se
Te
Po
.6 10.4 9.75 9.01
8.43
I.P.(eV) 13



Decreasesgradually
(e)
Electrogenativity : Electrogenativity decreases gradually.
O
S
Se
Te
Po
Electronegativity
3.5
2.5
2.4
2.1
2.0
(f) Melting and boiling points : The melting and boiling points increases gradually with increase in
atomic number.
(g) Allotropy:
Element
Allotropic forms
Oxygen
Ordinary oxygen and ozone
Sulphur
Rhombic, monoclinic, plastic ,Amorphous
Selenium
Red form (non-metallic), grey form (metallic form)
Tellurium
Crystalline and amorphous.
Polonium
 and  forms (Both are metallic forms)
9
(h)
(f)
Catenation: Oxygen and sulphur show the property of catenation.
Oxidation states: Oxygen being highly electronegative shows –2 oxidation state in its compounds,
since the electrogenativity decreases, the tendency to exhibit –2oxidation state decreases as we go
down in the group. However, positive oxidation states are exhibited by S, Se, Te and Po. In addition to
+2 oxidation state, +4 and +6 oxidation states are observed. This is due to the availability of d-orbitals
in these elements.
3.
Chemical characteristics:
(a) Hydrides: All the elements of this group form the hydrides of type H2M
M = O, S, Se, Te and Po, i.e., H2O, H2S, H2Se, H2Te and H2Po
(i)
Physical state: Water is colourless, odourless liquid while other hydrides are colourless, poisonous
gases with bad odours.
(ii)
Volatility: Water has low volatility (high boiling point) as hydrogen bonding brings association.
Volatility decreases from H2S to H2Te due to increase in molecular masses of the hydrides.
(iii) Covalent character: As the electronegativity difference between M and H decreases, the covalent
character of these hydrides increases from H2O to H2Te.
(iv) Thermal stability: The thermal stability decreases as the atomic mass increases. This is due to an
increase in M – H bond length.
(v)
Acidic nature: The acidic strength increases from H2O to H2Te.
(vi) Reducing nature: All hydrides except H2O act as reducing agents. The reducing nature increases as
the atomic number of the central atom increases. This is due to weakening. Of M – H bond as the
bond length increases with increase of size of M-atom.
(vii) All these hydrides are V-shaped. In these hydrides, the central atom is sp3 hydridized. The bond
angles are 104.5°, 92.5°, 91° 90 o in H2O, H2S, H2Se and H2Te, respectively.
(b)
Oxides: The most important oxides are of the type MO2 and MO3,SO2 and SeO2 are acidic oxides
and are soluble in water. TeO2 and PoO2 are insoluble in water. These are amphoteric oxides as they
dissolve in both acids and bases.
SO2 + H2O  H2SO3 (Sulphurous acid) SeO2 + H2O  H2SeO3 (Selenous acid )
TeO2 + 2NaOH  Na2TeO3 + H2O
(sodium tellurite)
2TeO2 + HNO3  2TeO2.HNO3 or Te2O3.(OH)NO3.
Trioxides: SO3, SeO3 and TeO3 are acidic in nature.
SO3 + H2O  H2SO4 (Sulphuric acid)
SeO3 + H2O  H2SeO4 (Selenic acid)
TeO3 + 3H2O  H6TeO6 (Telluric acid)
The acidic nature decreases on moving down the group
(c)
Oxyacids: S, Se and Te form similar oxyacids.
(d)
H2SO4
H2SeO3
H2TeO3
Sulphurous Acid Selenous acid Tellurous acid
Salts: Sulphites
Selenites
Tellurites
H2SO4
H2SeO4
H2TeO4
Sulphuric acid
Selenic acid
Telluric acid
Salts: sulphates
Selenates
Tellurates
Halides: S, Se and Te form hyxafluorides showing the maximum valency of six. They all involve
10
sp3d2 hybridization. Thus, hexafluorides possess octahedral structure. Many tetrahalides are known .
For Ex.SF4 is a gas, SeF4 is a liquid while TeF4 is a solid..
(e)
Oxyhalids: Only S and Se form oxyhalides. They are called thionyl and selenyl halides.
SOF2
SOCl2
SOBr2
SeOF2
SeOCl2
SeOBr2
These react with water readily.
SOCl2 + H2O  SO2 + 2HCl
In addition, sulhuryl halides are also known. SO2X2.
Anomalous Behaviour of Oxygen
The anomalous behaviour is attributed due to the following inherent characteristics.
(i)
Small size
(ii)
High electrogenativity
(iii) Non-availability of d-orbitals in the valency shell.
Sulphuric Acid (Oil of vitriol ) H2SO4.
Manufacture: Sulphuric acid is manufactured these days by the following two processes.
Lead Chamber Process:
Principle : The mixture containing sulphur dioxide, air and nitric oxide when treated with steam,
Sulphuric acid is formed.
2SO2 + O2 + 2H2O + [NO]  2H2SO4 + [NO]
(Air)
Catalyst
Catalyst
Nitric oxide acts as a catalyst in this reaction..
Contact Process:
Principle: The process involves the oxidation of sulphur dioxide by air in the presence of a catalyst.
Catalyst
2SO2
+
O2
2SO3
Sulphur trioxide is dissolved in 98% Sulphuric acid when oleum is formed.
H2SO4 + SO3  H2S2O7 (Oleum),
H2S2O7 + H2O  2H2SO4.
The Commonly used catalysts are platinum , ferric oxide or vanadium pentoxide. Nowadays, vanadium
pentoxide (V2O5) is preferred as it is cheaper and not poisoned by impurities. 13.
Oxyacids of sulphur : A large number of oxyacids are known in the case of sulphur either in free state or in
the form of salts or both. Oxyacids with S —S links are called thioacids.
1.
Sulphurous acid series :
(i)
Sulphurous acid, H2SO3
(ii)
Thiosulphurous acid, H2S2O2
(iii) Hyposulphurous acid, H2S2O4
(iv) Pyrosulphurous acid, H2S2O5.
(v)
Sulphuric acid, H2SO4
(vi) Thiosulphuric acid, H2S2O3
(vii) Pyrosulphuric acid, H2S2O7
(viii) Dithionic acid, H2S2O6
(ix) Polythionic acid, H2SnO6 (n = 3,4,5,6)
(x)
Peroxy monosulphuric acid , H2SO5 (Caro’s acid)
(xi) Peroxy disulphuric acid, H2S2O8 (Marshall’s acid)
11
The Elements of Group VIIA or 17 (Halogens)
17th or VIIA group of the periodic table (extended form ) consists of five elements; fluorine (F), chlorine (Cl),
bromine (Br), iodine (I) and astatine (At).
Physical characteristics
(a) Physical state: The tendency to form condensed molecules increases with increase in atomic number.
Fluorine and chlorine are gases at ordinary temperature , bromine is a highly fuming liquid while iodine is a
volatile solid.
F
Cl
Br
I
Gas Gas Fuming liquid Volatile solid
Halogens exist as diatomic covalent molecules. These molecules are held together by weak van der Waals’
forces. On account of weak intermolecular forces, the halogens are volatile in nature. With the increase in
size , these forces. Increase from fluorine to iodine and hence change in physical state occurs from gas (F 2
and Cl2) to solid (I2)
(b) Atomic and ionic radii: These increase from fluorine to iodine.
Element
F
Cl
Br
I
Atomic radiiÅ 0.72 0.99 1.14 1.33
Ionic raddiÅX 1.36 1.81 1.85 2.16
As we move from fluorine to astatine, an extra shell of electrons is added to each element. The addition of an
extra shell increases the atomic and ionic radii from fluorine to astatine.
(c) Atomic volume and density: These increase steadily from fluorine to iodine.
Element
F
Cl
Br
I
At. volume mL
17.1 18.7 23.5 25.7
Density in
1.108 1.1557 2.948 3.76
liquid state
(d) Energy and stability of X―X bond. With the increase of size , the bond length increases from fluorine
to iodine. Since the bond length of fluorine is minimum, its bond dissociation energy should be highest.
However, the bond dissociation energy of fluorine is less than Cl―Cl and Br―Br. Actually, the bond
dissociation energy should have decreased from fluorine to iodine but it starts decreasing from chlorine to
iodine.
X―X bond F―F Cl―Cl Br―Br
I―I
Bond lenth Å 1.42 1.99 2.28 2.67
Bond dissociation 38 57
45.5 35.6
Energy (kcal/mol)
The lower value of bond dissociation energy of fluorine is due to the high inter electronic repulsions between
non-bonding electrons in the 2p-orbitals of fluorine.
(e) Ionisation potential: The ionisation potentials of halogens are very high. This indication that the
halogens have a little tendency to loss electron to form X+ cation. However, the ionisation potentials
decrease from F to I, i.e., the tendency to form cation increases form F to I.
Element
F
Cl
Br
I
First ionisation
17.4 13.0 11.8 10.4
Potential (eV)
Decreases gradually
(f) Electron affinity : Electron affinity values are high in the case of halogens.
12
Element
F
Cl
Br
I
Electron affinity (eV) -3.6 -3.8
-3.5 -3.2
The low value of electron affinity of fluorine is probably due to small size of fluorine atom, i.e., electron
density is high which hinders the addtion of an extra electron.
(g) Electronegativity: The halogens have high values of electronegativity. Fluorine has the maximum
electronegativity. The electronegativity decreases from fluorine to idine.
Element
F
Cl
Br
I
Electronegativity
4.0 3.0 2.8 2.5
(h) Colour : Al the halogens are coloured.
F
Cl
Br
I
Light yellow, yellow green, Reddish brown deep violet
(I) Oxidation states: Since these elements have seven electrons in their valency shell, each element
tries to attain 8 electrons either by accepting an electron from an element which is less electronegative that
it or by sharing its unpaired electron in p-orbital with another element. When the halogen atom combines
with an element of lesser electronegativity, combines with an element hving higher electronegativity, it
exhibits +1 oxidation state. Fluorine, being most electronegative , always shows –1 oxidation state only.
These elements (i.e., Cl, Br and I) can show +3 , +5 and +7 oxidation states depending on the number of
singly occupied orbitals, i.e., oxidation states lie between –1 to +7.
ns np
nd
Halogen atom
(except F atom ) in1 unpaired electron , -1+1 oxidation
the ground state.
(ns2 sn5)
ns
np
nd
Halogen atom in
the first excited state
(ns2np4nd1) 3 unpaired electrons , +3 oxidation state
Halogen atom in
the second excited state ns
np
nd
(ns2 np3 nd2)
5 unpaired electron; +5 oxidation state
Halogen atom
ns
np
nd
in the third
excited state 7 unpaired electron;+7 oxidation state
(ns1 np3 nd3)
F atom has no p-orbitals in the valency shell, thus it cannot have any excited state and consequently
cannot show any of the higher oxidation states.
(j) Reduction potentials and oxidising nature: Standard reduction potentials of halogens are positive and
decrease from fluorine to iodine. Thus, halogens act as strong oxidising agents and their oxidising
power decreases from fluorine to iodine.
Element
F
Cl
Br
I
Reduction +2.87 +1.36 +1.06 +0.54
Potential
(Eo volt)
Decreases gradually
13
Oxidising nature also decreases
Fluorine is the strongest oxidising agent. It will oxidize other halide ions to halogens in solution or
when dry.
F2 + 2X-  2F- + X2
(X = Cl-, Br- or I-)
Similarly, chlorine will oxidize Br- and I- solution to Br2 and I2, respectively and Br2 will oxidize I- to I2.
in general, a halogen of low atomic number will oxidise the halide ion of higher atomic number.
The halide ions act as reducing agents. F- ion does not show any reducing nature but Cl-, Br- and Iion act as reducing agents and their reducing nature is in increasing order.
ClBrIReducing nature increases
Chemical characteristics
(a) Reactivity: Halogens are most reactive non-metallic elements, high reactivity is due to low
dissociation energies of the halogen molecules. Molecule N2
O2 H2
X2(halogen)
Dissociation 225 118.3 103.2
36.5 to 57.2
Energy (kcal/mol)
In halogens, fluorine is most reactive and iodine is least reactive.
(i)
Reaction towards water: Fluorine decomposes water very readily even at low temperature and in dark
forming mixture of O2 and O3, Cl2 decomposes water in presence of sunlight while bromine
decomposes water very slowly in presence of sunlight. Iodine does not decompose water.
 HX + HXO
X2 + H2O Sunlight
(Cl2 or Br2)
3F2 + 3H2O  6HF + O3
2F2 + 3H2O  4HF + O2
I2 + H2O  No reaction
(c) Oxides: The compounds of oxygen and fluorine are not called as oxides but oxygen fluorides as fluorine
is more electronegative than oxygen. The compounds of oxygen and rest of the halogens are termed
oxides. Halogens and oxygen do not combine directly with each other.
OF2 Cl2O Br2O I2O5
O2F2 ClO2 BrO2
Cl2O6 BrO3
Cl2O7 .
Oxides of chlorine are acidic and the acidic and the acidic nature increases as % of oxygen increases.
Thus, Cl2O is least acidic while Cl2O7 is most acidic oxide.
(d) Oxyacids: Fluorine does not form any oxyacid as it is more electronegative than oxygen. Other
halogens form oxyacids of the type HXO, HXO2, HXO3 and HXO4.
General Properties
(i)
The oxyacids are monobasic and form one series of salts.
(ii)
In oxyacids of the same halogen in different oxidation states, the acidic nature and thermal stability
increase while oxidising nature decreases with the increase of oxidation state of the halogen.
The acidic character and thermal stability are in the order:
HClO < HClO2 < HClO3 < HClO4
While the oxidising power is in the order:
HClO > HClO2 > HClO3 > HClO4
14
i.e, increasing oxygen content leads to:
(a) an increase in thermal stability
(b) an increase in strength of the acid.
(c) a decrease in strength of the acid.
(i)
In the oxyacids of different halogens in the same oxidation state, the thermal stability and acidic
nature, in general , decrease with increase of atomic number. However, in HXO3, HIO3 is most stable.
Halogen acids or Hydracids
All the halogens combine with hydrogen and form the covalent hydrides of HX type (X=F,CL,Br or I).
These hydrides are called hydracids or halogen acids.
Hydrogen fluoride (HF)
Hydrogen chloride (HCl)
Hydrogen bromide (HBr)
Hydrogen iodide (HI)
Properties
(i)
Physical state: Except H2F2, other hydrogen halides. They fume in air and have pungent odour.
These are colourless.
Their melting and boiling points increase with increase in atomic mass of the halogen. The low values
are due to covalent nature.
HCl
HBr
HI
o
Melting point C –111
-86
-50.8
o
Boiling point C
-85
-67
-35.5
o
H2F2 is a liquid with boiling point 19.5 C. this behaviour is due to association of HF molecules through
hydrogen bonding.
H―F-----H―F------H―F-----H―F-----(ii)
Stability : The bond strength H―X decreases from HF to HI. Thus, HF is most stable while HI is least
stable. The decrease in stability is due to decrease in electronegativity from F to I.
H―F H―Cl H―Br H―I
Dissociation
Energy (kcal mol-1) 136
105
86
70.
(iii) Acid strength:
(iv) Reducing nature: The reducing nature increases from HF to HI as the stability decreases from HF to
HI , HF does not show reducing nature.
Interhalogen Compounds
The halogens, on account of difference in their electronegavityties, combine with each other and form
binary covalent compounds, of Abn type which are called interhalogen compounds. A is always bigger
atom and B is smaller atom and n may have values 1,3,5 and 7 corresponding to oxidation states of
halogens. The value of n increases with increase of atomic number. These compounds are of four
types:
AB type
:
ClF, BrF, BrCl, ICl, Ibr
AB3 type
:
ClF3,BrF3, ICl3
AB5 type
:
BrF5, IF5
AB7 type
:
IF7.
General Properties :
15
(i)
Physical state :These may be gases (ClF, BeF,ClF3 , IF7) liquids (BrF3, BrF5, IF5) or solids
(ICI, IBr, IF3, ICI3 )
(ii)
Duanagbetuc batyre : All the electrons , bonding or non – bonding are present in pairs , the
interhalogen molecules are diamagnetic in nature .
(iii) Thermal stability : Thermal stability of AB type compounds increases with increase in
electronegativity difference between A and B, the more polar is the A
- B bond and hence greater is the thermal stability .
IF >BrF> ClF >ICI> IBr >BrCI
(iv) Reactivity
: AB type compounds are more reactive than A2 or B2 molecyles , since AB bind
is weaker than A-A and B-B bond . the order of reactivity of some inerhalogen compounds has been
found as :
ClF3> > BrF5 >If7 BrF3 > IF5> BrF
Polyhalides.
Halide ions often react with molecules of halogens or iterhalogen and from polyhalide ions. Iodine is only
slightly soluble in water. Its
solubility is greatly increased if some iodide ions are present in the solution.
The increase in solubility is due to formation of a polyhalide ion. I-3.
I- + I2  I-3.
Many polyhalides are known which contain two or three different halogens. For example K[ICl2], K[ICl4],
Cs[BrF] and K[IBrCl]. These are formed from interhalogens and metal halides.
Pseudohalides and Pseudohalogens
A few ion are known, consisting of two or more electronegative atoms of which at lest one is nitrogen,
that have properties similar to those of halide ions. These ions are called pseudohalide ions.
Pseudohalide ions are univalent and these form salts resembling halide salts. For example, sodium
salts are soluble in water but the silver salts are insoluble. The pseudohalide ions are:
Cyanide ion (CN-) Isocyanide ion (NC-)
Cyanate ion (OCN-) Fulminate ion (ONC-)
Thiocyanate ion (SCN-)
Isothiocyanate ion (NCS-)
Selenocyanate ion Tellurocyanate ion
(SeCN-)
(TeCN-)
Azide ion (N-3)
Azido carbon disulphide ion (SCSN3-)
16