General Characteristic
Valence configuration
ns2np2 (n= no of valence shell)
Common oxidation state
+2, +4 (due to inert pair effect) but +4 is stable
m+2 compounds are ionic and reducing agents m4+ compounds are covalent and oxidising agents
Bond energy and catenation property order
C-C> Si-Si> Ge-Ge> Sn-Sn
Stability of hydrides
CH4>SiH4>GeH4>SnH4>PbH4
Stability of Halides
CCl4>SiCl4>GeCl4>SnCl4>PbCl4>CF4>CCl4>CBr4>CI4
Acidic nature of dioxides
CO2>SiO2>GeO2>SnO2>PbO2
Oxides of Carbon and Silicon
Carbon Monoxide : (CO)
Neutral oxide
- Colorless and odourless
- Poisonous gas
Preparation:
(a) HCOOH−→−−−−−−−H2OConsH2SO4CO (b) H2C2O4−→−−−−−−−H2OCons⋅H2SO4CO (c) ZnO+C→ΔZn+CO Structure:
:C≡O sp hybridized
C−0 bond length=1⋅13A0 - Long pair on c atom is dented to certain metal to
form m←COcarbonyls. viz⋅NiCl4⋅ Fe(CO)5etc. - Carbon monoxide forms carboxy hemoglobin complex with blood which
may lead to fatal death because this complex is more stable than oxy-
hemoglobin complex.
- It is a good Reducing agent
ZnO+CO→Zn+CO2 Fe2O3+3CO→2Fe+3CO2
Carbon dioxide (CO2)
It is an acidic oxide
- Colorless and odourless gas
- It turns line water into milky
- Structure O=C=O+↔ O≡C−O− - C -atom sp hybridised bond order 2 M=0
C-O bond length =1⋅15A˙
Silicon dioxide (SiO2) (Silica)
Structure of SiO2 -Silica is acidic oxide
- Widely found as sand and Quartz
- Main forms of SiO2 are quartz, tridymite and crystobalite. - Colored quartz are used as gems
- Flint, opal, agate, any x and jaspel are amorphorus silica.
- Kieselglahr is siliceous rock composed if minute sea organisms.
- silica is soluble in HF
SiO2+4HF→SiF4→HFH2SiF6 - SiO2+2NaOH→Na2SiO3+H2O Water glass - Na2SiO3+2HCl→H2SiO3+2NaCl Silicic acid -SiO2 is a 3D polymer in which each Si is in sp3 hybridization. - SiC : carborundrum or Artificial diamond
SiO2+3C→ΔSiC+2CO
Organo-Silicon polymers
These are organo silicon polymers containing -Si-O-Si- linkages
and R2−SiO group as unit building block (R = methy or phenyl)
Silicates
Silicates
- Silicates make 95% of the Earth's crust, with silica and alumino -
silicate.
- They contain different mode of combination of (SiO4)−4 tetrahedral units.
- Si-O bond is 50% ionic 50% covalent.
Types of Silicates
(1) Ortho Silicates
(2) Pyrosilicates
(3) Ring (cyclic) Silicates
(4) Chain Silicates
(5) Sheet Silicates
(6) 3 D Silicates
Group 15 elements
Synopsis
Nitrogen, Phosphorus, Arsenic, Antimony and Bismuth belong to VA
group or 15th group of the periodic table.
The atomic numbers of N, P, As, Sb, and Bi are 7, 15, 33, 51 and 83
respectively.
Elements of Nitrogen family are called pnicogens and their compounds
are called pnictides.
The general valency shell electronic configuration of these elements
is ns2 np3.
Physical properties
Nitrogen is a diatomic gaseous molecule where as phosphorus is a
tetra-atomic solid. This is because nitrogen atoms are small in size
and can approach very close to one another so lateral overlap of p-
orbitals can takes place to form π -bonds.
Phosphorus atoms are larger in size hence lateral overlapping is not
possible. So, P4 molecules are formed by single bonds between P
atoms.
Nitrogen is chemically inert because N ≡ N energy is very high
(945.4 K.J. / mole).
Nitrogen and Phosphorus are non metals. Arsenic and Antimony are
metalloids.
Bismuth is a metal.
Atomic radius, metallic character, Density and
B.P. gradually increase from N to Bi.
The Nitrogen and Phosphorus are non conductors of heat and
electricity. Arsenic is a poor conductor Antimony and Bismuth are the
good conductors
of heat and electricity.
Ionisation potential, electronegativity, electron affinity gradually
decrease from N to Bi.
M.P. increases upto As and then decreases.
Physical properties of nitrogen
There is a considerable increase in covalent radius from N to P.
However, from As to Bi only a small increase in covalent radius is
observed.
This is due to the presence of competely filled dand /or forbitals
in heavier members.
Allotropes of nitrogen
Except bismuth all the elements of this group exhibit allotropy.
Nitrogen has two allotropes in the solid state. They are α -
Nitrogen (cubic crystalline) and β -Nitrogen (hexagonal
crystalline).
Phosphorus exists in a variety of forms. The most important forms of
phosphorus are white or yellow, red, α - Black, β - Black,
scarlet, violet.
White phosphorous is stored under water.
White phosphorous contains discrete P4 molecules.
In P4 molecule, the four P atoms are present at the corners of a
tetrahedron and bond angle is 60∘.
Most reactive form of phosphorus is white due to bond angle strain.
Phosphorous tetrahedron polymerises to form more inactive Red -
phosphorus.
Catenation capacity
The catenation capacity depends on bond energy Greater the bond
energy value, higher the catenation capacity
N - N B.E 83.7 K.J/mole
P - P B.E 79.06 K.J/mole
White phosphorus
1. It is a translucent white waxy solid. It is poisonous.
2. Insoluble water but soluble in carbon disulphide and glows in dark
( chemiluminescence).
3. It dissolves in boiling NaOH solution in an inert atmosphere given
PH3
P4+3NaOH+3H2O →PH3+3NaH2PO2 4. It readily catches fire in air to give dense white fumes of P4O10
Red Phosphorus
1. It is obtained by heating white phosphorus at 573K in an inert
atomsphere for several days. When red phosphorus is heated under
high
pressure, a series of phases of black phosphorus is formed.
2. Red phosphorus possesses iron grey lustre. It is odourless, non
poisonous and insoluble in water as well as in carbon disulphide.
3.Less reactive than white phosphorus.It does not glow in the dark.
Black Phosphorus
1. It has two forms α - black phosphorus and β -black phosphorus.
2. α - black is formed when red phosphorus is heated in a sealed
tube at 803K.
3. β -black is prepared by heating white phosphorus at 473K under
high pressure
Oxidation state
P similar to nitrogen exhibits all possible oxidation states between
+ III and +V in its hydrides, oxides
VA group elements exhibit -3, +3 and + 5 oxidation numbers.
Nitrogen exhibits all the oxidation states from -3 to +5.+ 5
oxidation state is unstable in Bi due to inert pair effect.
Stable oxidation number of Bi is +3 due to inert pair effect.
The stability of +5 state decreases down the group from N to Bi
In the case of nitrogen, all oxidation states from +1 to +4 tend to
disproportionate in acid solution. For ,
3HNO2 →HNO3+H2O+2NO Similarly, in case of phosphorus nearly all intermediate oxidation
states disproportionate into +5 and -3 both in alkali and acid
Compounds of pnicogens
Hydrides
The ability to donate lone pair (Lewis basic nature), stability,
solubility and basic strength of the hydrides decrease from NH3to
BiH3.
Reducing power of the hydrides increases from NH3to BiH3.
The ease of formation of these hydrides, their stability and their
tendency to from coordinate covalent bonds gradually decreases from
NH3 to BiH3
NH3 is best ligand and forms coordinate covalent bonds readily
MH3 type hydrides are trigonal pyramidal in shape.
In NH3 molecule central atom will make use of SP3 hybrid orbitals.
In MH3 type hydrides, the bond angle decreases from NH3 to BiH3 due to
incre
Properties of Hydrides
Except NH3 other hydrides have little or no tendency to form
coordinate covalent bonds (to donate e−pair)
Though N has greater EN ammonia is the strongest electron donor of
all the hydrides of VA group elements. This is
(1) Because the small size of the nitrogen atom (because of small
size
e− density is more on sp3 hybrid orbital compared to p and other elements.)
(2) in other hydrides greater M-H bond length leads to weakening of
the covalent bond.
(3) The lone pair of e− is spread over a larger atom. As a result of this e− density on the atom and e− donating nature (basic nature) decreases.
Stability order to hydrides
NH3 >> PH3 >> AsH3 >> SbH3 BiH3
Order of basic nature
NH3 > PH3 > AsH3 > SbH3 > BiH3
NH3 forms hydrogen bonds with water
From N to Bi. E.N decreases and so that polarity of M-H bond
decreases hence their solubility also decreases
Trends in some properties of Hydrides of
V A group elements
1. M.Ps PH3 < AsH3 < SbH3 < NH3
2. B.Ps PH3 < AsH3 < NH3 < SbH3
3. B.Ls NH3 < PH3 < AsH3 < SbH3
4. B.Es NH3 > PH3 > SbH3 > AsH3
5. B.As NH3 > PH3 > AsH3 > SbH3
As pure p orbitals of As and Sb are involved, theHMH bond angle in
AsH3and SbH3 would be expected 90∘. But due to repulsions between M-H
bonds, the angle increases to 91∘. 481
Ammonia is only a mild reducing agent while BiH3 is the strongest
reducing agent amongest all the hydrides.
Preparation of phosphine
Phosphine is prepared by the reaction of calcium phosphide with water
or dilute HCl
Ca3P2+6H2O→3Ca(OH)2+2PH3 Ca3P2+6HCl →3CaCl2+2PH3 In the laboratory, it is prepared by heating white phosphorus with
concentrate NaOH solution in an inert atmosphere of CO2 P4+3NaOH+3H2O →PH3+3NaH3PO2
Reactions of phosphine
It is slightly souble 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, the
corresponding phosphides are obtained.
3CuSO4+2PH3 →Cu3P2+2H2SO4 3HgCl2+2PH3 →Hg2P2+6HCl Phosphine is weakly basic and like ammonia, gives phosphonium
compounds with acids e.g., PH3+HBr →PH4Br When pure it is non inflammable, but becomes inflammable owing to the
presence of P2H2 or P4 vapours.
To purify if from the impurities, it is absorbed in HI to form
phosphonium iodide (PH4 I) which on treating with KOH gives off
phosphine.
PH4I+KOH →KI+H2O+PH3
Oxides-Synopsis
These elements form the series of oxides -Trioxides (M2O3) and
Pentoxides (M2O5).
Nitrogen forms number of oxides due to Pπ Pπ multiple bonding
between N and oxygen atoms.
As oxidation number of the element increases, acidic nature of its
oxides increases.
Acidic nature of pentoxides is more than that of trioxides
Oxides of nitrogen
N2O : Nitrous Oxide (or) Nitrogen monoxide.
It is also known as laughing gas.
It is prepared by heating ammonium nitrate.
It is a colourless neutral oxide.
It is a linear molecule.
The structure of N2O is
Usually, N2O is administered to the patient to put him sleep
:N≡N˙→O¨..:↔N..=N˙=O¨:
NH4NO3−→−−HeatN2O+2H2O N - N - O
113 pm 119pm
Linear
NO : Nitric oxide
It is formed as an intermediate in the manufacture of HNO3 by
catalytic oxidation of NH3 in presence of Pt.
4NH3+5O2−→ΔPt4NO+6H2O NO formed during lightening stage
It is a colourless, neutral gas.
It is paramagnetic due to the presence of unpaired electron (11
valency electrons).
It gives reddish brown gases in air.
It readily reacts with O2 as
2NO+O2→2NO2 (Reddish brown gas) Its structure is :N=...O: 2NaNO2+2FeSO4+3H2SO4 →Fe2(SO4)3+2NaHsO4+2H2O+2NO
N2O3 : Nitrogen trioxide : It is also known as Nitrogen sesqui oxide.
It is formed by cooling an equimolar mixture of NO and NO2.
It is a blue liquid at low temperature.
It is an acidic oxide.
It is anhydride of Nitrous acid.
NO2 (or) N2O4 : Nitrogen dioxide or Dinitrogen tetroxide. It is obtained by heating Lead Nitrate
2Pb(NO3)2−→Δ2PbO+4NO2+O2 It is a reddish brown poisonous gas soluble in water.
It becomes a colourless solid on cooling due to the formation of
dimer N2O4. It dissolves in water giving HNO2 and HNO3. So it is called mixed
anhydride.
NO2 is an odd electron molecule and exhibits paramagnetic property. In dimeric state (N2O4) it is colourless and diamagnetic in nature.
N2O5 : Dinitrogen pentoxide. It is obtained by dehydrating HNO3 with P2O5.
It is the anhydride of Nitric acid.
It is a powerful oxidising agent.
It is a colourless solid.
It disolves in water to give nitric acid.
N2O5+H2O →2HNO3
Oxides of phosphorus
P4O6: Phosphorus trioxide.
It is obtained by burning phosphorus in limited supply of air.
It is the anhydride of phosphorus acid.
It dissolves in cold water to form phosphorus acid.
In P4O6 each phosphorus is surrounded by three oxygen atoms.
It is an acidic oxide.
Number of P-O-P bonds are six
P4O10: Phosphorus pentoxide.
It is obtained by burning phosphorus in excess of air or oxygen.
It is the anhydride of phosphoric acid.
It dissolves in water to form H3PO4.
In P4O10 each phosphorus is surrounded by four
oxygen atoms.
Number of P-O-P bonds are six
It is a strong dehydrating agent.
From N2O3 to Bi2O3 acidic nature decreases and
basic nature increases
Acidic nature decreases or basic nature increases from N2O5 to Sb4O10
P4O10+6H2O →4H3PO4
Halides
VA group elements form trihalides of the type MX3 and pentahalides of
the type MX5.
NF3 does not undergo hydrolysis.
NCl3 on hydrolysis gives NH3 and Hypochlorous
acid.
PF3 is weakly reactive to water.
NCl3+3H2O →NH3+3HOCl PCl3 on hydrolysis gives HCl and H3PO3. PCl3+3H2O →H3PO3+3HCl
Tri halides are covalent.
Trihalides use the sp3 hybridised orbitals of the central atom.
In the formation of PCl5, the central phosphorus will make use of
SP3 d hybrid orbitals.
Trihalides have trigonal pyramid structure.
Nitrogen cannot form NCl5 because it has no d-orbitals in the valency
shell.
Bi does not form BiCl5 due to inert pair effect.
Penta halides use the sp3 d hybridised orbitals of the central atom.
Pentahalides have trigonal bipyramidal structure.
The extent of hydrolysis decreases from NX3 to BiX3.
In these hydrolysis reactions, the non metallic nature decreases or
metallic nature increases from N to Bi
Phosphorus pentachloride
1. It is prepared by the reaction of white phosphorus with excess of
dry chlorine
P4+10Cl2 →4PCl5 2. It can also be prepared by the action of SO2Cl2 on phosphorus. P4+10SO2Cl2 →4PCl5+10SO2
Properties
PCl5 is yellowish white powder and in moist air. It hydrolyses to POCl5 and finally gets conveted to phosphoric acid PCl5+H2O →POCl3+2HCl POCl3+3H3O→H3PO4+3HCl When heated it sublimes, but decomposes on stronger
heating PCl5−→−−HeatPCl3+Cl2 It reacts with organic compounds containing -OH group converting them
to chloro derivaties.
C2H5OH+PCl5→C2H5Cl+OCl+HCl Finely divided metals on heating with PCl5 give corresponding chlorides.
2Ag+PCl5→2AgCl+PCl3
Sn+2PCl5→SnCl4+2PCl3
In the solid state it exists as an ionic solid.
[PCl4]+[PCl6]− in which the cation, [PCl4]+ is tetrahedral and the anion. [PCl4]+ octahedral.
Nitrous acid
Nitrous acid (HNO2) :
Nitrous acid is unstable except in dilute solutions In the laboratory
it is prepared by the addition of ice cold dilute acid to Barium
nitrite
Ba(NO2)2+H2SO4 →BaSO4+2HNO2 (ice cold)
Its solution is slightly bluish in colour due to the presence N2O3
On standing it undergoes auto oxidation-reduction in acidic solution
3HNO2 →HNO3 +2NO +H2O In this reaction,
In HNO2 →HNO3 O.S. of 'N' changes from +3 to +5 In HNO2 →NO O.S. of N changes from +3 to +2 i.e HNO2 as oxidant changes to NO and as reductant changes to HNO3"
At low temperatures HNO2 reacts with aromatic primary amines and
gives diazonium compounds Diazonium compounds can be converted into
different substituted aromatic compounds.
Structure of nitrous acid
STRUCTURE OF (HNO2) :
HNO2 exists in two tautomeric forms i.e in two structural isomers.
Nitric acid(Aqua fortis)
The structure of nitric acid is
It is a very strong oxidising agent. It oxides non-metals to their
corresponding oxides or oxoacids
P4+20HNO3 →4H3PO4+20NO2+4H2O C+4HNO3 →CO2+4NO2+2H2O
This is a monobasic acid.
It is a strong oxidising agent.
The molecular of pernitric acid is HNO4.
Reactions of nitric acid
Concentrated nitric acid is a strong oxidising
Agent and attacks most metals excpet noble metals such as gold and
platinum
3Cu+8HNO3(dilute)→ 3Cu(NO3)2+2NO+4H2O Cu+4HNO3(conc.) → Cu(NO3)2+2NO2+2H2O
Zinc reacts with diltue nitric acid to give N2O and with concentrated acid to give NO2
4Zn+10HNO2(dilute)→ 4Zn(NO3)2+5H2O+N2O Zn+4HNO3(conc.) → Zn(NO3)2+2H2O+2NH2
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.
i) It is oxidised to iodine to iodic acid
I2+10HNO3 →2HIO3+10NO2+4H2O (ii) Carbon to carbon dioxide,
C+4HNO3→CO2+2H2O+4NO2 (iii) Sulphur to H2SO4 S8+48HNO3→8H2SO4+48NO2+16H2O (iv) Phosphorus to phosphoric acid
P4+20HNO3 →4H3PO4 +20NO2+4H2O
Hypophosphorus acid (H2PO2)
It is prepared by the heating yellow or white p with dilute Ba(OH)2
6H2O+2P4+3Ba(OH)2 →3Ba(H2PO2)2+PH3↑ from Ba(H2PO2)2, H3PO2 is obtained by hydrolysis.
1) H3PO2 in monobasic acid and a very strong reducing agent is basic solutions and it is oxidised to H3PO3 2) Meta phosphorous acid (HPO2) It is mono basic acid normally exist as a cyclic compound due to
polymerisation.
3) Ortho phosphorous acid (H3PO3) It is prepared by disolving P4O6 in cold H2O P4O6+6H2O →4H3PO3 or P(OH)3
Ortho phosphoric acid
1. It is prepared by dissolving P4O10 in water P4O10+6H2O→4H3PO4 2. It is a weak tribasic acid and has oxidising properties.
3. It forms three type of salts
4. Solid H3PO4 absorbs water and forms a colourless syrupy liquid (syrupy phosphoric acid)
Ortho phosphoric acid is prepared in the lab by the action of HNO3 on phosphorus.
H3PO4 is manufactured by heating bone ash or phosphorite rock with dil. H2SO4. H3PO4 is a tribasic acid. H3PO4 forms three types of salts. Primary phosphates : H2PO−4 Secondary phosphates :HPO2−4 Tertiary phosphates : PO3−4 In H3PO4 phosphorous atom is sp3 hybridised.
Metaphosphoric acid (HPO3)
Meta phosphoric acid:It is f orm ed by heating pyrophosphoric acid or
orthophosphoric acid to
870kH3PO4−→−−−H2O520kH2P2O7−→−−−H2O870kHPO3 It is a transparent glassy solid.
HPO3 is a monobasic acid and its salts are called meta phosphates.
Important features of oxyacids of phosphorus
In all these oxyacids, phosphorous is tetra hedrally surrounded by
atoms (generally)
In all these oxyacids, at least one OH group is linked to the
phosphorous atoms. The hydrogen atoms in OH groups are ionisable, and
responsible for the acidic nature.
P-H bonds are responsible for reducing properties of the acids
phosphoric series of acids do not have P-H bonds.
1. All oxoacids contain at least one P=O and one P-OH bond.
2. Hypophosphorous acid is a good reducing agent as it contains two
P-H bonds and reduces, for , AgNO3 to metallic silver.
4AgNO3+2H2O+H3PO2 → 4Ag+4HNO3+H3PO4
Preparation of HNO3
HNO3 is prepared on large scale by
1) Birkland-Eydes process (Arc process)
2) Ostwalds process (from ammonia)
Birkeland Edye process
Used at places where electric power is cheap
Principle : -
N2+O2−→−−−−−−Electric arc2NO; ΔH=180.7kJ 2NO+O2 →2NO2 4NO2+O2+2H2O→4HNO3
Ostwald's process
NH3 mixed with air in 1 : 7 or 1 : 8 when passed over a hot platinum gauze catalyst is oxidised (95%) to NO
4NH3+4O2 −→−−−1155Kpt gauze4NO+6H2O+1275KJ The liberated heat keeps the catalyst hot.
The NO gas is cooled and mixed with oxygen to get NO2 in oxidation chamber.
Then it is passed into warm water under pressure in presence of
excess air where HNO3 is formed. 4NO2+O2+2H2→4HNO3 The acid formed is about 61% concentrated.
Uses of nitric acid
In the manufacture of fertilisers like basic calcium nitrate
[CaO.Ca(NO3)2]
In the preparation of explosiv es like TNT, nitroglycerine etc. as
nitration mixture along with H2SO4
In the preparation of perfumes, dyes and medicines
HNO3 is a very strong oxidising agent used in the oxidation of cyclohexanol or Cyclohexanone to adipic acid.
p-xylene to terepthalic acid
In the preparation of artificial silk i.e cellulose nitrate"
Preparation of ammonia
Ammonia is manufactured by following process.
1) From coal 2) By Habers process
3) by Cyanamide process
Ammonium salt on heating with an alkali gives ammonia gas.
NH4Cl+NaOH →NaCl+NH3+H2O 2NH4Cl+Ca(OH)2→CaCl2+2NH3+2H2O Nitrolim is a mixture of calcium cyanamide and graphite (CaCN2 + C)
Calcium cyanamide on hydrolysis gives ammonia gas.
CaCN2+3H2O →CaCO3+2NH3
Preparation of ammonia-Haber's process
On large scale, ammonia is prepared by Haber's process.
In Haber's process, ammonia is synthesized directly from elements.
The nitrogen and hydrogen used in the Haber's process must be very
pure
N2+3H2⇔2NH3 : ΔH=−93.63KJ Conditions : Temperature : 725 to 775 K
Pressure : 200- 300atm
Catalyst : Finely divided iron
Promoter : Molybdenum or Ox ides of Potassium and Aluminium
Nitrogen and Hydrogen are mixed in the ratio 1:3
Dehydrating agents like P2O5, Con. H2SO4, anhydrous CaCl2 are not used for drying NH3, because they react with ammonia.
Ammonia is dried over CaO (Quick lime).
Reactions of ammonia and ammonium
Ammonia gives brown precipitate with Nessler's reagent K2[HgI4]. The of the precipitate formed in the above, i.e. Hg2O.NH2I(Iodide of millions base).
Nessler's reagent is a mixture of KI, HgCl2 and NaOH.
Ammonia forms ammonium salts with acids, e.g.,
NH4Cl,(NH4)2SO4 , etc. As a weak base, it precipitates the hydroxides (hydrated oxides in case of some metals) of man metals from their
salt solutions. For ,
ZnSO4(aq)+2NH4OH(aq)→ Zn(OH)2(s)+(NH4)2SO4(aq)(Whiteppt) FeCl3(aq)+NH4OH(aq) → Fe2O3.xH2O(s)+NH4Cl(aq)(brownppt) Due to Lewis basic nature, it forms complex compounds with metals
like Cu2+,Ag+ Cu2+(aq)+4NH3(aq)↔[Cu(NH3)4]2+(aq) (blue) (deep blue)
Preparation of dinitrogen
In the laboratory, dinitrogen is prepared by treating an aqueous
solution of ammonium chloride with sodium nitrite
NH4Cl(aq)+NaNO2(aq) → N2(g)+2H2O(l)+NaCl(aq)
Small amounts of NO and HNO3 are also formed in this reaction; these impurities can be removed by passing the gas through aqueous
sulphuric acid containing potassium dichromate. It can also be
obtained by the thermal decomposition of ammonium dichromate.
(NH4)2Cr2O7−→−−HeatN2+4H2O+Cr2O3
Very pure nitrogen can be obtained by the thermal decomposition of
sodium or barium azide.
Ba(N3)2 →Ba+3B2 2NaN3→2Na+3N2
Properties of dinitrogen
Dinitrogen is a colourless, odourless, tasteless and non-toxic gas.
At higher temperatures, it directly combines with some metals to form
predominantly ionic nitrides and with non-metals, covalent nitrides.
A few
typical reactions are.
6Li+N2−→−−Heat2Li3N 3Mg+N2−→−−HeatMg3N2
Calcium superphosphate
Calcium super phosphate is Ca(H2PO4)2+2(CaSO4.2H2O) Calcium super phosphate is a mixture of calcium dihydrogen phosphate
and gypsum.
It is a phosphatic fertilizer.
It is soluble in water.
Uses
Uses of phosphine
The spontaneous combustion of phosphine is technically used in Holmes
signals.(Mixture of calcium carbide and calcium phosphide.
It is also used in smoke screens.
Uses of ammonia
As a refrigerant
As a solvent
In the manufacture of Ammonium sulphate, Urea and other fertilizers.
In the manufacture of HNO3 by ostwalds process
Uses of dinitrogen
Liquid dinitrogen is used as a refrigerant to preserve biological
materials, food items and in cryosurgery.
Uses of bleaching powder
Bleaching Powder is used as Bleaching agent in textile and paper
industry.
Bleaching Powder is used for the sterilization of drinking water.
Percentage of Chlorine in bleaching powder is 56%
Bleaching Powder is used for the manufacture of chloroform.
It is oxidising agent and chlorinating agent
Group 17 elements
Synopsis
The elements Fluorine, Chlorine, Bromine, Iodine and Astatine are
present in 17th column of periodic table .
The halogens have ns2 np5 electronic configuration in their outermost
shell.
Astatine is a synthetic radio active element. It is called Radio
active Halogen.
All the Halogens have a strong tendency to take up one electron to
attain stable noble gas configuration. Hence, they are very reactive
and occur only in the combined state.
Halogens react among themselves forming interhalogen compounds, which
are more reactive than the halogen molecules.
Oxidation state
Fluorine always exhibits a fixed oxidation state of -1 in its
compounds because it is the most electro negative element
Chlorine, Bromine and Iodine show both negative and positive
oxidation states
Chlorine, Bromine and Iodine show -1, +1, +3, +5 and +7 oxidation
states. Higher oxidation states are due to the presence of vacant d-
orbitals.
Chlorine, Bromine and Iodine form 1, 3, 5 and 7 bonds due to the
presence of vacant d-orbitals.
Shape and Hybridisation halogen compounds
Chlorine, Bromine and Iodine
contain only one unpaired electron in ground state. They can show -1
or +1 oxidation state in ground state.
In ground state, halogen atom contain one unpaired electron and three
lone pair of electrons.
The shape of XA type molecule is Linear.
(X = less electronegative halogen,
A =more electronegative halogen)
Cl, Br and I in their first excited state contain 3 unpaired
electrons and form 3 bonds and exhibit +3 oxidation state.
The possible hybridisation of Cl, Br and I atoms in first excited
state is sp3d.
In the first excited state, halogen atom contain 3 unpaired electrons
and 2 lone pairs.
The shape of XA3 type of molecules is T.
(X = Cl, Br, I)
Cl, Br and I in their second excited state contain 5 unpaired
electrons and exhibit +5 oxidation state forming 5 bonds.
The possible hybridisation of Cl, Br and I in second excited state is
sp3d2.
In the second excited state halogen atom contain 5 unpaired electrons
and 1 lone pair of electrons.
The shape of XA5 type of molecules is square pyramid
(X = Cl, Br, I)
Cl, Br and I in their third excited state contain 7 unpaired
electrons and exhibit +7 oxidation state forming 7 bonds.
The possible hybridisation in halogen atom in third excited state is
sp3d3.
The shape of XA7 type of compounds given by halogens is pentagonal
bipyramid.
Preparation of fluorine
Fluorine is prepared by Whytlaw Grays method
The products of electrolysis of fused KHF2 are hydrogen at cathode
and flourine at anode.
Fluorine prepared in the electrolytic cell is passed through U-tubes
containing sodium fluoride to remove HF vapours present in Fluorine
as NaHF2
In Whytlaw Grays method, rectangular copper vessel acts as cathode
and a graphite rod acts as anode.
In Whytlaw Grays method, graphite anode is surrounded by a perforated
copper diaphragm to avoid mixing up of H2
and F2.
Abnormal behavior of fluorine
Abnormal behaviour of fluorine is due to
a) small size
b) highest electronegativity
c) low dissociation energy for F-F bond and
d) 2 electrons only in the penultimate shell while other halogens
have 8 electrons.
The abnormal characteristics of flourine are
a) F2 exhibits only -I oxidation state
b) In its hydride it forms hydrogen bonding and forms HF2
ion but of other halogens hydrides do not show hydrogen bonding .
c) It combines directly with carbon while others do not, even under
drastic conditions.
d)F2 has a lower E.A compared to Cl2 even thorugh
F2 is the most electronegative element.
e) Fluorides have maximum ionic character.
Fluorine is oxidising agent.
2KHSO4+F2 →K2S2O8 H2S+4F2 →2HF+SF6 Glass dissolves in HF only due to the formation of Hydro fluoro
silicic acid (H2SiF6).
SiO2+4HF →2H2O+SiF4
SiF4+2HF →H2SiF6 HF is used for etching or marking glass.
Fluoro Chloro Carbon is called Freon. It is used as a refrigerant.
Polymeric tetra fluoro ethylene is called Teflon. It is used as an
anti corrosive plastic.
Fluorine is used in the seperation of U235 and U238 in the form of
UF6gases based on atmolysis.
NaF and Na3AlF6 are used as insecticides.
DDFT is used as fungicide.
F2 is used in rocket fuels
Preparation of chlorine
Chlorine can be prepared by the oxidation of HCl
with MnO24HCl+MnO2→ MnCl2+2H2O+Cl2
Chlorine is prepared when a mixture of common salt and concentrated
H2SO4 is used in place of HCl.
4NaCl+MnO2+4H2SO4 →MnCl2+ 4NaHSO4+2H2O+Cl2
By the reaction of HCl on potassium permanganate.
2KMnO4+16HCl→2KCl+2MnCl2+8H2O+5Cl2
Deacon's process and Nelson's cell-
Deacons process : By oxidation of hydrogen chloride gas by atmosphere
oxygen in the presence of CuCl2 at 723 K.
4HCl+O2−→−−−CuCl22Cl2+2H2O
In Nelson's cell method, Chlorine is manufactured by the electrolysis
of Brine or an aqueous solution of sodium chloride.
In Nelson's cell, a perforated steel vessel acts as cathode and
graphite rod acts as anode.
A perforated steel cathode is used in Nelson's cell to prevent the
mixing up of Cl2 and NaOH
In Nelson's cell the product at anode is Cl2 and the products at
cathode is H2, NaOH.
In Nelson's cell for the manufacture of Cl2 the valuable by-products
are NaOH and H2.
Reactions of chlorine
Chlorine reacts with dry slaked lime to form bleaching powder
Ca(OH)2+Cl2→CaOCl2+H2O Chlorine form addition compounds with SO2, CO and NO.
SO2+Cl2→SO2Cl2 CO+Cl2→COCl2 It has great affinity for hydrogen. It reacts with compounds
containing hydrogen to form HCl.
H2+Cl2 →2HCl H2S+Cl→2HCl+S C10H16+8Cl2 →16HCl+10C Clorine is oxydising agent.
H2S+Cl2 →2HCl+S Na2SO3+H2O+Cl →Na2SO4+HCl Na2S2O3 +Cl2+H2O →Na2SO4+S+2HCl Chlorine is used as a bleaching agent in paper and textile industry.
Chlorine is used for the sterilization of drinking water.
It is used in the extraction of metals like gold and platinum.
Chlorine water on standing loses its yellow colour due to the
formation of HCl and HOCl. HOCl gives nascent oxygen which is
responsible for oxidising and bleaching properties of chlorine.
Cl2+H2O →2HCl+O Coloured substabce + O → Colourless substance
Compounds of chlorine
COCl2 is called phosgene. It is poisonous gas.
CCl3. NO2 is called tear gas.
Cl−C2H4−S−C2H4−Clor(C2H4Cl)2 S is called Mustard gas. It is used as war gas.
Dichloro diphenyl trichloro ethane is known as DDT. It is a
fungicide.
Physical properties of halogens
Physical state
Halogens exist as diatomic covalent molecules.
The only type of attractions between Halogen molecules are
vanderwaals forces.
Fluorine and Chlorine are gases. Bromine is a liquid and Iodine is
solid at room temperature.
Iodine exhibits sublimation property because of the presence of weak
vanderwaals forces.
The atomic and ionic radii gradually increases from Fluorine to
Iodine.
Atomic volumes of Halogens increase from Fluorine to Iodine.
All halogens are coloured. By absorbing different quanta of
raidiation they display different colour
F2 - Yellow Cl2 - Greenish yellow Br2 - Red I2 - Violet
Density
The Densities of Halogens increases from Fluorine to Iodine
Melting and Boiling points
The Melting and boiling points increase gradually from fluorine to
iodine, due to increase in vander waals forces.
The NonMetallic Nature decreases from fluorine to iodine.
Ionization potential
The Ionisation potentials of Halogens are very high.
The Ionisation potentials decrease from Fluorine to Iodine, due to
the increase in atomic size.
Iodine having the lowest value of Ionisation potential has some
tendency to form I + ion.
Iodine is the only Halogen that can form both positive and negative
charge ions.
Iodine is called reducing Halogen..
Electron affinity and Electronegativity Form-short
Electron affinity values of halogens are very high.
The electron affinity of fluorine is less than chlorine though it is
most electronegative. This is due to its small size. Repulsions
between newly added electron and the electrons already present in its
small 2p orbital
are high. Chlorine high electron density in a relatively compact 2p -
subshell
Electron affinity values of Halogens are in the order.
Cl > F > Br > I
Electro negativity values of Halogens are very high.
Electro negativity values of Halogens decreases from Fluorine to
Iodine.
Element with highest electro negativity is Fluorine.
Element with highest electron affinity is chlorine.
Bond dissociation energy
Bond dissociation energies of Halogens are in the order Cl2 > Br2 >
F2 > I2.
According to Mulliken, the high bond dissociation energies
of Cl2,Br2,and I2are due to multiple bonds formed by d - p combinations or overlapping.
The low bond dissociation energy of F2 is due to absence of d- p combination as it does not possess d-orbitals.
According to Coulson,the low bond dissociation energy of fluorine is
due to more repulsion between the lone pairs of electrons on the two
smaller fluorine atoms and Inter-nuclear repulsions
Flourine is extremely reactive due to very low bond dissociation
energy. Hence, it is called super halogen
Solubility
Halogens are soluble in water which follow the order
F2 > Cl2 > Br2 > I2
Halogens being non-polar do not dissolve to a significant extent in a
polar solvent like water.
In CCl4,CS2 or paraffins, Cl2, Br2 and I2 gives yellow, brown and violet colour respectively.
The solubility of iodine in water is enhanced in presence of KI
KI+I2⇌KI3⇌ K+ +I−3
Reactivity
Halogens are highly reactive elements they can react with metals as
well as non-metal and other substances. the order of reactivity of
Halogens is
F2 >> Cl2 > Br2 > I2
Reactions of halogens
Reaction with water
Halogens are sparingly soluble in water because they are non-polar
covalent molecules. The solubility of Halogens decrease from F2 to I2
Fluorine decomposes water to liberate a gaseous mixture of (O2+O3) 2H2O+2F2 →4HF+O2 3H2O+3F2 →6HF+O3
Chlorine reacts with water to form HCl and HOCl
Cl2+H2O →HCl+ HOCl Chlorine water contains HCl and HOCl
Chlorine acts as a bleaching agent in the presence of water or
moisture due to formation of HOCl.
The bleaching action of chlorine in the presence of water or moisture
is due to oxidation or liberation of nascent oxygen.
HOCl → HCl + (O)
The reaction of iodine with water is non-spontaneous. In fact, it can
be oxidised by oxygen in acidic medium just the reverse of the
reaction observed with flourine
4I−(aq)+4H+(aq)+O2(g) →2I2(g)+2H2O(1)
Reactivity towards oxygen
Halogens form many oxides with oxygen but most of them are unstable.
Fluorine forms two oxides OF2 and O2F2 .
Reactivity towards halogens
OF2andO2F2 Both are strong fluorinating agents. O2F2 oxidises plutonium to PuF6 and the reaction is used in removing plutonium as PuF6 from spend nuclear fuel.
Chlorine, bromine and iodine form oxides in which the oxidation
states of these halogens range from +1 to +7. A combination of
kinetic and thermodynamic factors lead to the generally decreasing
order of stability of oxides formed by halogens, I > Cl > Br .
The higher oxides of halogens tend to be more stable than the lower
ones.
Chlorine oxides, Cl2O,ClO2,Cl2O6 and Cl2O7 ae highly reactive oxidising agents and tend to explode.
Reaction with hydrogen
All the Halogens directly combine with Hydrogen to form Hydrides
(a) H2+F2−→−23K2HF It is a fast reaction and takes place even in the dark and is highly energetic
(b) H2+Cl2 −→−−−−Sunlight2HCl it is Slow in dark but fast in Sunlight
(c) H2+Br2−→Δ2HBr It does not take place at room temperature. Takes place at 593 K in
Sunlight.
(d) H2+I2⇌2HI It takes place in the presence of Pt as catalyst at 713 K and is a
reversible change
Acidic strength of halides
The stability of the hydrides decreases from HF to HI due to decrease
in their dissociation energies.
The stability order of hydrogen halides is HF > HCl > HBr > HI
The order of acidic strengths of halides- HF < HCl < HBr < HI.
The stability of halides decrease due to increase of bond
dissociation energy.
The order of bond dissociation energy- HF > HCl > HBr > HI
Reaction of chlorine with NH3
When excess chlorine reacts with ammonia to form an unstable Nitrogen
trichloride and HCl.
3Cl2+NH3→3HCl+NCl3
Chlorine reacts with excess ammonia to give NH4Cl liberating
Nitrogen.
3Cl2+8NH3→6NH4Cl+N2
Reactions with alkalies
Fluorine reacts with cold and dil. NaOH to form NaF, H2O & OF2.
2F2+2NaOH →2NaF+H2O+OF2
Fluorine reacts with hot and conc. NaOH liberating oxygen gas
2F2+4NaOH →4NaF+2H2O+O2 Cl2,Br2 and I2 react with cold and dil. NaOH to form halide and hypo halites. The oxidation number of halogen changes from 0 to -1& +1
Cl2+2NaOH →NaCl+NaOCl+H2O
Cl2,Br2 and I2 react with hot and conc. NaOH to form halide and halates. The oxidation state of halogen changes from 0 to -1 and + 5.
3Cl2+6NaOH →5NaCl+NaClO3+3H2O
Ionic character of halides
The order of the ionic character of the halides
MF > MCl > MBr > MI where M is a monovalent metal. Halides in higher
oxidation state will be more covalent than the one in lower
oxidation state.
Oxidising power
Halogens are strong oxidising agents.
Fluorine is the strongest oxidising agent eventhough chlorine has
maximum electron affinity.
The magnitude of the enthalpy change in the reaction, when halogen
changes to a hydrated ion can be estimated by the application
of BORN-HABER cycle.
The overall change in enthalpy
(ΔH)=[ΔH1+ΔH2+D2−E−ΔH3] ΔH1= Enthalpy of fusion, ΔH2= Enthalpy of vapourisation ΔH3= Enthalpy of hydration D= Enthalpy of Dissociation E= Electron affinity
Due to low heat of dissociation of F2 molecule and high hydration
energy of F− ion, fluorine acts as strong oxidising agent.
A Halogen with lower atomic number oxidises a Halide ion of higher
atomic number.
Chlorine oxidises Bromides to Bromine and Iodides to Iodine.
Cl2+2KBr →Br2+2KCl Cl2+2KI →I2+2KCl
Bromine oxidises iodides to Iodine
Br2+2KI →I2+2KBr
Oxyacids of chlorine
Hypochlorous acid HClO or HOCl+1
Chlorous acid HClO2 + 3
Chloric acid HClO3 + 5
Percloric acid HClO4 + 7
Cl - O bond length decreases from OCl− to ClO−4
Cl - O bond energy increases from OCl− to ClO−4 except for ClO−3
Hypochlorous Acid
Chlorine atom in ClO− ion is sp3 hybridised with lone pair electrons.
ClO− ion is stable due to strong tendency to form Ppi - dpi bonding
between filled p-orbitals of oxygen and vacant d-orbitals of
chlorine.
Between one oxygen atom and chlorine atom there is σ bond
Chlorous Acid
Chlorine atom in ClO−2 ion is sp3 hybridised with two lone electron
pairs.
The shape of ClO−2 ion is angular
ClO−2 ion contains 2σ and one π bonds.
The bond angle is 111∘
Chloric Acid
Chlorine atom in ClO3 ion is sp3 hybridised with one lone electron
pair.
The shape of ClO3 ion is pyramidal.
ClO−3 ion contains 3σ and 2 π bonds.
In ClO−3 ion O-Cl-O bond angle is 106θ
Perchloric Acid
Chlorine atom in ClO−4 ion is sp3 hybridised with no lone pair
electrons.
The shape of ClO−4 ion is tetrahedral.
ClO−4 ion contains 4σ and 3π bonds.
The O-Cl-O bond angle is 109∘ 28`.
Perchloric acid is dimerized due to hydrogen bond
Compounds of chlorine
Bleaching powder-Synopsis
of Bleaching Powder is CaOCl2.
Bleaching Powder is also called chloride of lime.
The chemical name of Bleaching Powder is calcium chloro hypo chlorite.
The oxidation states of chlorine in Bleaching Powder are -1 and + 1.
Bleaching Powder is prepared by the action of dry chlorine and dry
slaked lime. The process called Bachmann process.
Properties of Bleaching powder
Bleaching Powder is unstable. On long standing it decomposes to form
CaCl2 and Ca(ClO3)2
6CaOCl2→5CaCl2+Ca(ClO3)2 The cold aqueous solution of Bleaching Powder contains Ca2+, Cl− and
OCl− ions.
The hot aqueous solution of Bleaching Powder contains Ca2+, Cl− -and
ClO−3 -ions.
Reactions of bleaching powder
Bleaching Powder reacts with excess dil. acids to liberate chlorine
gas. The amount of chlorine liberated is called Available Chlorine.
CaOCl2+H2SO4→CaSO4+H2O+Cl2 Similar reaction takes place when CO2 is passed over bleaching powder
paste prepared with H2O
CaOCl2+CO2→CaCO3+Cl2↑
A good sample of Bleaching Powder contains 35 to 38% of available
chlorine.
Interhalogen compounds
Interhalogen compounds
Halogens react with each other to produce a number of INTERHALOGEN
COMPOUNDS XXn where n = 1 , 3, 5 or 7.
The stability of the interhalogen compounds increases as the sixe of
the central atom increases.
CIF3 and BrF3 are used for the production of UF6 in the enrichment of
uranium (U235).
Interhalohen compounds can be prepared by direct combination or by
the action of halogen on lower interhalogen compounds.
Cl2+F2 −→−−437K2CIF (equal volume)
Cl2+3F2 −→−−573K2CIF3 (excess)
Br2+3F2 →2BrF3 (excess)
Inter halogen
compound Hybridisation
O.N
bonds σbonds Shapes
XA No
hybridization +1 1 Linear
XA3 sp3d +3 3 T-shape
XA5 sp3d2 +5 5 Square Pyramidal
XA7 sp3d3 +7 7 Pentagonal
bipyramidal
Interghalogen compounds are covalent and diamagnetic. ClF is gas and
the rest are solids or liquids at 298 K.
Being polar interhalogen compounds are more reactive then halogens
except fluorine.
All interhalogen compounds undergoes hydrolyses giving halideion
.Derived from the smaller halogen and hypohalite [when XX'] halite
when [when XX3'3] halate [when XX'5] perhalate [when XX'7].
Anion derived from the longer halogen.
XX′+H2O →HX+HOX
Group 16 elements
Synopsis
Oxygen, Sulphur, Selenium, Tellurium and Polonium are the elements of
VIA group or 16 vertical column of the periodic table.
The first four elements are collectively known as chalcogens(ore
forming elements), since many metals occur as oxides (or) sulphides
in nature.
Poloniumis a radio active element and shortlived.
The elements belong to p-block since the differentiating electron
enters the p-orbital.
The outer electronic configuration of these elements is ns2np4 Oxyzen is the most abundant element. It constitutes 46.6% of earths
crust,21% of air and 89.1% of ocean by weight
Properties
As Atomic weight increases, the formation of multiple bond between
identical atom s decreases.
Atomic radius increases from O to Te.
Ionisation energy and Electro-negativity decreases from O to Po.
Density increases from oxygen to polonium.
Melting points and boiling points changes irregularly as follows. O
<S<Se<Te>Po.
The large difference in the melting point of Oxygen and Sulphur is due
to a change from diatomic O2 to octa-atomic S8 which increases
magnitude of Van der waals forces.
Electron affinity values decrease from S to Te. Electron affinity of
O< S due to small size of Oxygen atom.
Order of electron affinity : S>Se>Te>O (E.A1)
Oxygen exists as a diatomic molecule (O−2), Sulphur, Selenium, Tellurium exist as octa-atomic molecules. (S8,
SeS8, TeS8).
Metallic character increases from O to Po.
Oxygen, Sulphur are strongly non-metallic.
Selenium and Tellurium are semimetals and are considered to be
metalloids.
Polonium is a metal.
Oxidation state
The common oxidation state of these elements is -2, because they
have s2p4 configuration in their outer most orbit. Oxygen shows positive oxidation states in fluorides.
Oxidation state of oxygen in O2F2 is +1and in OF2 is +2.
The Oxidation state of oxygen in a peroxide is -1 and in a super
oxide is frac−12 Oxygen cannot exhibit oxidation number beyond +2, due to absence of
vacant d-orbitals.
The other elements exhibit -2, +2, +4, +6 oxidation states due to
availability of vacant d orbitals in its valency shell.
Allotropes of oxygen
All VI A group elements exhibit allotropism except Te
Oxygen occurs in two non metallic forms (a)Oxygen, O2 (b)Ozone, O3
Oxygen is paramagnetic as it contains two unpaired electrons (as per
Molecular Orbital Theory).
Ozone is a triatomic diamagnetic allotropic form of oxygen. It is
unstable and decomposes to O2,2O3→3O2
Allotropes of sulphur
Sulphur has more number of allotropic forms. All these are non-
metallic.
Allotropes of Sulphur are :
a. α - Sulphur or Rhombic sulphur. .
b. β -Sulphur or monoclinic sulphur or prismatic
sulphur.
c. γ - Monoclinic sulphur. .
d. x - Sulphur or plastic sulphur. . α,β and γ forms of sulphur are crystalline in
nature and possess puckered ring structures (S8 )
(crown configuration)
Catenation capacity
Oxygen and Sulphur show catenation tendency.
Catenation is maximum in sulphur upto 10 atoms
(H2Sn)(n=2−10) In case of Oxygen it is limited to 2 atoms only .
(H2O2)
Anomalous behavior of oxygen
Due to small size and high electronegativity, oxygen exhibits
anomalous behaviour
Ex: oxygen forms hydrogen bonds in H2O but sulphur cant form in H2S The absence of d - orbitals in oxygen limits its covalency to four.
Hydrides
The binary compounds of VI A group elements with hydrogen are called
hydrides.
VIA group elements can form the hydrides of the H2 M (M=VIA group element).
All these hydrides are covalent compounds.
H2O - hydrogen oxide (water) H2S - hydrogen sulphide H2Se - hydrogen selenide H2Te - hydrogen telluride H2Po - hydrogen polonide
The thermal stability Order :
H2O > H2S > H2Se > H2Te M-H bond length in the hydrides increases from
H2O to H2Po The aqueous solutions of these hydrides behave as weak acids. The
acidic strength increases from
H2O to H2Te. Order : H2O < H2S <H2Se <H2Te Ka (dissociation constant) of aqueous solutions of these hydrides
increases from H2O to H2Te The reducing property increases from H2O to H2Po Covalent character Order : H2O >H2S > H2Se >H2Te In the formation of H2O, Oxygen atom is sp3 hybridised. In H2S and other hydrides, pure p-orbitals are involved in bonding. At room temperature, water is a liquid and the other hydrides are
colourless, foul smelling, toxic gases.
When compared to other hydrides, water has abnormal high boiling
point. This is due to intermolecular hydrogen bonding in water.
Boiling point order: H2 <H2Se<H2Te<H2O Order of volatile nature is H2S>H2Se>H2Te>H2O
The molecules have bent structures (VSEPR theory).
The bond angle decreases from H2O to H2Po (H2O−104∘28′;H2S−92∘30′;H2Se−91∘;H2Te−90∘;H2Po−90∘)
Oxides
Oxides are two types:
1. Simple Oxide (Na2O,Al2O3 etc..) 2. Mixed Oxides
Mixed oxides : - Formed by the combination of two simple oxides eg :
Red lead, PbO4
(PbPO2.2PbO),Fe3O4(FeO+Fe2O3) Simple oxides classified as
( i ) Acidic oxides - oxides of non metals which give acids when
dissolved in water are called acidic oxides
eg. CO2,NO2,P2O5,SO2,SO3,Cl2O7 etc... CO2+,H2O →,H2CO3 carbonic acid SO2+H2O →H2SO3 (Sulphurous acid) The metalic oxides of high oxidation states
eg Mn2O7,V2O5 and CrO3 (ii) Basic oxides
(a) lonic oxide. Oxides of alkali and alkaline earth metals
eg Na2O,CaO,BaO. In water they give basic solutions.
Na2O+H2O →2NaOH CaO+H2O →Ca(OH)2 (b) Covalent oxides-oxides of transition metals are covalent in
nature eg CuO, FeO. Insoluble in water
(iii) Amphoteric oxides - The oxides which react with acids and
alkalies are known as amphoteric oxides eg ZnO,Al2O3,SnO3 etc Al2O3(s)+6NaOH(aq)+3H2O(l)→2Na3[Al(OH)6](aq) (iv) Neutral oxides - such oxides do not combine with an acid or a
base eg : NO,N2O,CO,H2O etc.
Compounds of sulphur
Properties of SO2
SO2 is a colourless gas with pungent smell.
SO2 can easily be condensed to a liquid ( at room temperature by
using pressure of two atm) which is used as a non-aqueous solvent.
Liquid SO2 is highly soluble in water and forms
Sulphurous acid.( SO2 H2O ⇔ H2 SO3 )
SO2 acts as a Lewis base due to the presence of lone-pair of
electrons.
It acts as mild reducing agent in acidic solutions and a strong
reducing agent in basic solutions.
Chemical properties of sulphur
SO2 reduces acidified K2Cr2O7 into Cr(III) sulphate.
K2Cr2O7 +H2SO4+3SO2 →
K2SO4+Cr2(SO4)3+H2O
Its reacts readily with NaOH solution forming Na2SO3 . Which then
reacts with more SO2 to give sodium hydrogen sulphite
Na2SO3+H2O+SO2→2NaHSO3
Moist SO2 reacts with Fe (+III) salts to give Fe (+II) salts. It
decolourises acidified KMnO4 (VII) Solutions
2Fe+3+SO2+2H2O → 2Fe+2+SO−24+4H+ 5SO2+2MnO−4+2H2O →5SO−24+4H++2Mn+2
SO2 is a bleaching agent. It bleaches the vegetable colouring matter
by reducing. In this process it is oxidised to H2SO4. This bleaching
process is temporary.
SO2+2H2O →H2SO4+2(H) Coloured matter + 2(H) → colourless matter.
Reaction of SO2 with Cl2 :
SO2+Cl2−→−−−−sunlighthvSO2Cl2 (Sulphurly Chloride)
(The reaction is also takes place in presence of charcoal)
Structure of SO2
In SO2, sulphur atom is sp2 hybridised. It is an angular molecule
with a lone pair.
The bond angle of O-S-O is 119∘ . 30!
In SO2 2 σ bonds and 2 π bonds
[(onedπ−pπ) & (onepπ−pπ)] are present. SO2 molecule has to resonance structures.
Preparation and properties of SO3
SO3 is prepared by reaction of SO2 with O2 in the
presence of the catalyst platinum or V2O5 or NO+NO2
2SO2+O2⇔2SO3
SO3 is the anhydride of sulphuric acid or sulphuric anhydride.
H2SO4→SO3+H2O
Commercially it is not possible to react SO3 directly with H2O. Hence SO3 dissolved in conc. H2SO4 to give oleum (H2S2O7). Then it is dissolved in water to get H2SO4. SO3+conc.H2SO4→H2S2O7(oleum) H2S2O7+H2O→2H2SO4 Oleum is also called fuming sulphuric acid or pyro sulphuric acid
Structure of SO3
In gaseous SO3, the central atom sulphur undergoes sp2 hybridisation.
The shape of SO3 molecule is planar triangular and O-S-O bond angle
is 120∘.
S-O bond length : 143pm or 1.43A.
In SO3 3 σ bonds and 3 π bonds are present they
are [(two dπ−p π) &(one pπ−pπ)]bonds.
SF6
Sulphur hexa fluoride is formed by the direct combination of sulphur
and fluorine.
S+3F2 −→ΔSF6 SF6 is a colourless, odourless, non-flammable gas.
SF6 is highly stable and inert compound due to steric reasons.
SF6 is a covalent compound and has low boiling point.
In SF6, sulphur atom is sp3d2 hybridised.
All F-S-F angles are 90∘. The shape of SF6 molecule is octahedral
SF4.SCl4
SF4 is also prepared by reaction between Cobalt trifluoride with
Sulphur.
S+4CoF3→SF4(g)+4CoF2 SF4 is thermally more stable than lower fluorides.
SF4 is highly reactive gas and a good fluorinating agent.
SCl4 can be prepared by the direct reaction between sulphur and
chlorine.
S+2Cl2→SCl4
SF4 and SCl4 are Lewis acids since they can accept lone pairs of
electrons readily to form hexahalides using halide ions.
SF4 and SCl4 can act as Lewis bases by donating lone pairs of
electrons.
In SF4 and SCl4, sulphur atom undergoes sp3d hybridisation ,
Tetra halide (SF4, SCl4 ) molecules have trigonal bipyramidal
structure with one corner of equatorial position occupied by a lone
pair of electrons(sea-saw structure)
S2Cl2
Sulphur on heating with chlorine gives S2Cl2. This on saturation with
chlorine gives SCl2.
2S+Cl2→S2Cl2 S2Cl2+Cl2 →2SCl2
Sulphur dichloride reacts with ethylene to form di (2-chloro ethyl)
sulphide, commonly known as mustard gas
SCl2+2CH2=CH2→S(CH2−CH2−Cl)2
In SCl2, the sulphur atom is in sp3 hybridisation, with two positions
occupied by lone pairs.
In SCl2, the lone pairs distort the tetrahedral angle from 109∘281 to
103∘
Oxyacids of sulphur
The oxyacids of sulphur are classified into four series.
a) Sulphurous acid series
b) Sulphuric acid series
c) Thionic acid series
d) peroxy acid series
The hybridisation of S in all oxyacids is sp3.
Oxyacids of Sulphur with S - S linkage are called thioacids.
Salt of Caro's acid is called permonosulphate and salt of Marshall's
acid is perdisulphate or persulphate.
Distillation of H2S2O8 with water gives H2SO5 which on further
hydrolysis gives H2O2.
Basicity of all oxo-acids of Sulphur is 2.
Hybridization of sulphur in
SO−23,SO2−4,S2O−23,S2O−24 ions is sp3
Preparation of sodium thiosulphate
Hydrated Sodium Thiosulphate (Na2S2O3.5H2O) is called Hypo.
When Alkaline or neutral Na2 SO3 solution is boiled with flowers of
sulphur gives hypo.
Na2SO3 + S → Na2S2O3 (solution) (excess) (hypo)
Hypo is prepared by oxidation of sodium sulphide or sodium poly
sulphide with air
2Na2S5+3O2 −→Δ2Na2S2O3+6S
Chemical properties of sodium thiosulphate-hypo
Hypo with a dilute solution of AgNO3 gives a white precipitate which
changes to yellow, brown and finally black due to the formation of
Ag2S.
With concentrated solution of Hypo, AgNO3 gives no precipitate. This
is because silver thiosulphate (a white ppt) formed in the reaction
is easily soluble in excess of Hypo forming a complex,
Na3[Ag(S2O3)2]
(sodium argentothiosulphate).
Silver halides dissolve in hypo solution to give sodium
argentothiosulphate.
2Na2S2O3+AgBr →Na3[Ag(S2O3)2]+NaBr
This reaction is made use of in photography. This is known as fixing
in photography.
Hypo removes excess chlorine in moist condition.
So hypo is used as an antichlor.
Na2S2O3+H2O+Cl2 →Na2SO4+S+2HCl
Hypo reduces Iodine to form sodium tetrathionate. In Na2S4O6,
oxidation number of Sulphur is +2.5
2Na2S2O3+I2 →Na2S4O6+2Nal
Synopsis- H2SO4
Because of its wide applications in industry, it is called King of
chemicals. It was also called as OIL OF VITRIOL.
There are two important methods of manufacturing sulphuric acid.
1) Lead chamber process
2) Contact process
Contact process
The steps involved are :
i) Burning of sulphur (or) sulphide ores (like iron pyrites) in air
to get SO2
S +O2 →SO2 4FeS2+11O2 →2Fe2O3+8SO2
(ii) Conversion of SO2 to SO3 catalytically 2SO2+O2−→−−−Δcatalyst2SO3
(iii) SO3 is absorbed in 98% H2SO4 to get oleum SO3+H2SO4 →H2S2O7
Oleum is diluted with water to get sulphuric acid of desired
concentration
H2S2O7+H2O →2H2SO4
The key step in the process is catalytic oxidation of SO2 with O2 to give SO3 in presence of catalyst V2O5 Forward reaction is : Exothermic and Δn = -ve According to Le Chatlier's principle to favour forward process the
following conditions are to be maintained.
I. High pressure is preferred. But actually 2 atm pressure is
maintained.
II. Low temperatures are preferred.
Physical properties of H2SO4
During dillution,conc.acid is
slowly added to water as acid dissolves in
H2O liberates large amount of heat
When pure H2 SO4 is cooled with ice it solidifies to colourless
crystals which melts at 10.380C.
High boiling point and high viscosity of H2SO4 is due to the fact
that H2SO4 molecules are associated together by H-bonding.
Chemical properties of H2SO4
Its chemical reactions are due
to
i) Low volatility,
ii) Strong acidic charecter
iii) Strong affinity for water.
iv) Ability to act as oxidising agent
It ionises in water in two steps as
H2SO4(aq)+H2O(l)→H3O+(aq)+HSO−4(aq) Ka1= very high HSO−4(aq)+H2O(l) →H3O+(aq)+SO−24(aq) Ka2= is very less (1.2×10−2)
It is very good dehydrating agent.
It removes water from corbohydrates as
C12H22O11 →12C++11H32O C6H12O6 →6C++6H2O
Hot conc. H2SO4 is moderetly strong oxidising agent. (Strength is in between H3PO4 and HNO3) eg : -
Cu+2H2SO4(conc)→CuSO4+SO2+2H2O C+2H2SO4(conc) →CO2+2SO2+2H2O
Compounds of Oxygen
OF2
VIA group elements can form mono halides of the type M2X2;
dihalides (MX2); tetrahalides (MX4) and hexa halides
(MX6).
Since the electro-negativity of fluorine is greater than oxygen,
the compounds of fluorine and oxygen are called fluorides of oxygen.
OF2 is pale yellow gas. It is prepared by passing F2 gas through a
very dil. solution of NaOH.
2NaOH+2F2→2NaF+H2O+OF2
Structure of OF2 is angular.
FOF bond angle in OF2 is 103∘.
O2F2
O2F2 is prepared by passing silent electric discharge through a mixture of F2 and O2 at very low temperature.
F2O2 →O2F2 O2F2 has open book like structure.
In O2F2 molecule bond angle is 109∘ 31!
dihedral angle is 87∘ 30!
Preparation method of oxygen gas
By heating oxygen containing salts such as chlorates, nitrates &
permanganates
2KClO3−→−−−MnO2Δ2KCl+2O2 2NaNO3−→Δ2NaNO2+O2 2KMnO4−→ΔK2MnO4+MnO2+O2
Thermal decomposition of oxides of metals which are in lower part of
electrochemical series
2Ag2O(s)→4Ag(s)+O2(g) 2HgO(S)→2Hg(l)+O2(g) 2Pb2O4(S)→6PbO(S)+O2(g)
Chemical properties of oxygen gas
Oxygen directly reacts with nearly all metals and non-metals except
some metals like Au, Pt and noble gases.
With metals :- 2Ca + O2 → 2CaO
4Al + 3O2 Al2 O3
With non metals:- P4 + 5O 2 → P4O10
C + O2 → CO2
With other compounds
ZnS + 2O2 → 2ZnO + 2SO2
CH4 + 2O2 → CO2 + 2H2O
Some compounds are catalytically oxidised.
2SO2+O2−→−−V2O52SO3
Preparation of ozone
Ozone is prepared by subjecting cold,
dry oxygen gas to the action of silent electric discharge.
Formation of ozone is an endothermic, reversible reaction.
Physical properties of ozone
O3 is a pale blue,pungent smelling poisonous gas,dark blue liquid,
violet black solid.
Ozone is thermodynamically unstable.
Decomposition is associated with increase in volume.
In the decomposition, heat liberates and the entropy increases(Δ S
is positive) for the decomposition of ozone in to oxygen ΔG value is
negative.
It is heavier than air and is only slightly soluble in water.
It is highly soluble in turpentine oil, glacial acetic acid, or
carbontetrachloride
It decolourises organic colouring matter by oxidation.
Structure of ozone
Ozone is an angular molecule with a bond angle of 116∘ 49
O-O bond length is 1.28 A0. Ozone has two resonating structures. It is a diamagnetic molecule.
Uses
Uses of ozone
It is used as germicide and disinfectant.
It is used for sterilizing water.
It is used in improving the quality of atmosphere at crowded places
(tube railways, mines, cinema halls etc.,).
It is used for bleaching oils, oil paintings, ivory articles.
It is used in the manufacture of artificial silk and synthetic
camphor.
It is used to identify the unsaturation in carbon compounds.
A mixture of O3 and C2 N2 is known as (cyanogen) and is used as Rocket
fuel.
Uses of hypo-Sodium thiosulphate
Hypo is used as a
1. fixing agent in photography
2. antichlor in textile industry
3. volumetric reagent to estimate iodine in volumetric analysis.
4. antiseptic in medicine.
Uses of H2SO4
It is extensively used in
a) Petroleum refining
b) Manufacture of paints, dye stuffs
c) Detergent industry
d) Storage batteries (Lead storage batteries)
e) Manufacture of nitrocellulose products
f) Pickling agent
g) Laboratory reagent
Group 18 elements
Synopsis
Helium, neon, argon, krypton, xenon and radon are known as noble
gases. Their symbols are He, Ne, Ar, Kr, Xe and Rn.
The atomic numbers of He, Ne, Ar, Kr, Xe, and Rn are 2, 10, 18, 36,
54 and 86 respectively.
The first compound of noble gas was prepared by N. Bartlett. The
compound is xenon hexafluoro platinate (IV) Xe[PtF6]
Synopsis
Valence shell electron configuration of helium is 1s2 (duplet) Valence shell electron configuration of a noble gases
is ns2np6 (except He)
Noble gas atoms (except helium) have 8 electrons in their valence
shell. This type of electron arrangement is known as octet.
Noble gas atoms are chemically inert. So, they are also known as
inert gases.
The inertness of noble gases in chemical reactions is attributed to
the octet structure they have in their valence shell.
Noble gases are present only in extremely small concentrations in the
air. So, they are also known as rare gases.
Noble gases are isolated from the air, so they are also known
asaerogens.
Earlier the valency (i.e., combining capacity) of noble gases is
thought to be zero. So, noble gases are branded as Zero group
elements in the periodic table
Physical Properties
Oxidation state of noble gases is zero.
All Noble gases are monoatomic due to value of specific heat ratio
(Cp/Cv) 1.66.
They all display a regular gradation in their physical properties
such as At.Wt., B.P. Freezing point, density are increases from He to
Rn.
Noble gases are slightly soluble in water, however solubility
increases down the group.
Adsorption by charcoal: Extent of adsorption increases down the group.He < Ne < Ar < Kr < Xe
Liquefaction of gases: It is difficult to liquify noble gases due to
week Vander Walls forces however ease of liquification increases down
the group due to increase of Vander Walls forces.
Helium can form interstitial compounds with metals. In such
compounds, the host inert gases (Kr, Xe) are trapped into the
cavities of crystal lattices of host (organic or inorganic
compounds) these are non stoichiometric in nature.
eg: Quinol clathrate Xe.6H2O In clathrates, the bonding between noble gas atom and water is dipole
- Induced dipole interaction.
Properties of noble gases
Atomic number, atomic weight, radius of atom, density increases.
Van der Waals forces of attraction increases. So, boiling point
increases from He to Xe.
Heat of vapourisation, solubility in water increases.
Ionisation potential decreases
The electron affinity of inert gas is nearly equal to zero
Uses of noble gases
Noble gases are used to provide inert atmosphere in the extraction of
metals like Mg, Ti etc., and welding works which involve metals like
Mg, Al etc.
Helium is used as a heat transfer agent in nuclear reactors.
A mixture of 80 % helium and 20 % oxygen by volume is used by deep sea
divers for respiration.
He + O2 mixture is used to provide relief for the asthma patients in
their respiratory problems.
Liquid helium is used as a cryogenic liquid, to provide low
temperature.
Helium is used in gas thermometers and in electrical transformers.
Helium is used to fill the tyres of big aeroplanes because it is
lighter than air.
Neon glow lamps are used as signal lights, and as beacon lights for
safe air navigation.
Argon is used in filling electrical bulbs.
Kr -85 is used to measure thickness of metal sheets and joints.
Kr -85 is used in electronic tubes for voltage regulations.
Xenon is used in photographic flash bulbs.
Radon is used in making the ointments used in the treatment of cancer.
Radon is used as a substitute for x-rays in industrial radiography.
Isolation of noble gases
Ramsay-Rayleigh First method
CO2 present in pure and dry air is removed by passing it over soda lime and potash solutions.
O2 present in pure and dry air is removed by passing it over red hot copper
2Cu+O2→2CuO N2 present in pure and dry air is removed by passing it over red hot
magnesium metal 3Mg+N2→Mg3N2 The residual air obtained by removing CO2,O2,N2 from pure and dry air is a mixture of noble gases.
Ramsay Rayleigh Second Method
A mixture of pure and dry air, oxygen in 9:11 ratio by volumes is
subjected to electrical discharge applying a potential difference
of 6000- 8000 volts between the platinum electrodes.
Oxides of nitrogen are formed due to the following reactions
N2O2→2NO 2NO+O2→2NO2 The oxides of nitrogen formed are absorbed in aqueous NaOH solution.
2NO2+2NaOH→NaNO3+NaNO3+H2O
The residual air obtained by removing N2 and O2 in the form of oxides of nitrogen in a mixture of noble gases, having traces of oxygen .
The traces of oxygen present in the mixture of noble gases is removed
by absorbing it in alkaline pyrogallol solution.
Fischer Ringe's process
A mixture of 90% calcium carbide and 10% anhydrous calcium chloride
is heated to 1073K.
Pure and dry air is passed over this mixture.
Nitrogen and oxygen are removed from the air due to the following
reactions.
CaC2+N2→CaCN2+C C+O2→CO2 CO2+C→2CO
The CO is removed by passing it over red hot cupric oxide.
CuO+CO→Cu+CO2
The CO2 is removed by absorbing it in potash solution. CO2+2KOH→K2CO3+H2O
The mixture of noble gases is dried first over conc.H2SO4 and then over P4O10
Dewar's method
It is used to separate the mixture of noble gases. This method
depends on the adsorption of the noble gases on activated charcoal in
Dewar's flask
Except helium all the noble gases can be adsorbed over activated
coconut charcoal.
Adsorption of noble gases increases with an increase in the atomic
weight.
Noble gas with low atomic weight can be adsorbed on the coconut
charcoal at low temperature.
Noble gas with high atomic weight is adsorbed on the coconut charcoal
only at high temperature
The noble gas adsorbed in the charcoal comes out when the charcoal is
heated.
Compounds of noble gases
Synopsis
Xenon forms a number of compounds with fluorine and with oxygen.
Helium and neon cannot form compounds because they have no excited
state. Krypton forms a limited number of compounds. Eg: KrF2,KrF4
Xenon fluorides
Xenon forms three types of fluorides. They are XeF2, XeF4 and XeF6
In XeF2, oxidation state of xenon is +2 and hybridisation of xenon
atom is sp3d
The shape of XeF2 is linear with a bond angle 180∘ and bond length 2
A0.
XeF2 molecule has 3 lone pairs (equatorial postions) and 2 bond pairs
of electrons.
In XeF4, oxidation state of xenon is +4 and hybridisation of xenon
atom is sp3d2.
XeF4 molecule has 2 lone pairs (at axial position) and 4 bond pairs
of electrons.
The shape of XeF4 is square planar with a bond angle 90∘ and bond
length 1.95 A0.
In XeF6, the oxidation state of xenon is +6 and hybridisation of
xenon atom is sp3d3.
XeF6 molecule has 1 lone pair and 6 bond pairs of electrons.
The shape of XeF6 is distorted octahedron (or) Pentagonal bipyramid
(or) Capped octahedron.
The bond angle values in XeF6 are 144∘ and 90∘
Xenon oxides
Xenon forms two oxides XeO3 and XeO4.
XeO3 is unstable, decomposes to form Xe and O2
XeO3 is a colourless & hygroscopic substance with explosive nature.
In XeO3, oxidation state of xenon is +6 and hybridisation of xenon is
sp3.
The shape of XeO3 is pyramidal and bond angle is 103∘, due to the
presence of a lone pair of electrons on Xe.
XeO3 has 3 sigma and 3 pi bonds.
In XeO4 oxidation state of xenon is +8 and hybridisation of xenon is
sp3
The shape of XeO4 is tetrahedron and bond angle is 109∘ 28'
XeO4 has 4 sigma and 4 pi bonds.
Halogens
Characteristics of Halogen compounds
Valence electronic configuration- ns2np5 Oxidation states- (-)1, +1, +3, +5, and +4 and +6 states are possible
in oxides, oxy acids.
F exhibits only -1 oxidation state.
Electro negativity order- F > Cl > Br > I
Electro affinity order- Cl > F > Br > I
Oxidizing nature order- F > Cl > Br > I
Reducing nature order- Cl−<Br−<I− Bond Dissociation Energy order- Cl-Cl > Br-Br > F-F > I-I
Formation of hydrogen halides
2KHF2→Δ2KF+H2F2 or
CaF2+H2SO4→ΔCaSO4+H2F2 H2+Cl2→2HCl or
2NaCl+H2SO4→Na2SO4+2HCl H2+Br2→Pt2HBr or
2KBr+H3PO4→K3PO4+3HBr 3KI+H3PO4→K3PO4+3HI
Action with halogen
2HCl+F2→2HF+Cl2 2HBr+Cl2→2HCl+Br2 2HI+Br2→2HBr+I2
Trends in periodic properties
Stability order- HF > HCl > HBr > HI
Acidic strength- HI > HBr > HCl > HF
Reducing Nature- HF < HCl < HBr < HI
Boiling point order- HCl < HBr < HI < HF
Dipole moment order- HI < HBr < HCl< HF
Action with AgNO3
HCl+AgNO3→HNO3+AgCl↓(white ppt)
HBr+AgNO3→HNO3+AgBr↓ (Pale yellow ppt)
HI+AgNO3→HNO3+AgI↓ (yellow ppt)
Cl2O
Cl2O (Dichlorine monoxide) Oxidation state +1
Cl2+2HgO→orCl2O+HgO⋅HgCl2 Cl2+H2O⇌2HClO (golden yellow solution)
Chlorine dioxide (ClO2)
Oxidation state = +4
2AgClO2+Cl2→2ClO2+2AgCl 2NaClO2+Cl2→2ClO2+2NaCl 2ClO2+H2O→HClO2+HClO3
Dichlorine (Cl2O6) Hex oxide
Oxidation state = +6
It is mixed anhydride of chloric and per-chloric acid.
2ClO2+O2→0⋅cCl2O6 Cl2O6⇌2ClO3 (Diamagnetic) (Paramagnetic)
Cl2O6+H2O→HClO3+HClO4
Dichlorine hepatoxide (Cl2O7)
Oxidation state = +7
2HClO4+P2O5→Cl2O7+2HPO3 Cl2O7+H2O→2HClO4
Hypo chlorous acid (HClO)
Cl−O−H+ (linear) Cl=sp3 2Cl2+2HgO+H2O→2HClO+HgO⋅HgCl2 2CaO⋅Cl2+H2O+CO2→2HClO+CaCl2+CaCO3 It is unstable and decomposes
HClO→HCl+(O)
Chlorous acid (HClO2)
Ba(ClO2)2+H2SO4→2HClO2+BaSO4 It under goes auto oxidation.
2HClO2→HClO+HClO3
Chloric acid (HClO3)
Ba(ClO3)2+H2SO4→2HClO3+BaSO4 6NaOH+3Cl2→not5NaCl+NaClO3+3H2O
Per chloric acid (HClO4)
2KClO4+2H2SO4→2HClO4+2KHSO4 It is the strongest acid and fames in moist air.
2HClO4+Mg→Mg(ClO4)2+H2
Acidic and Oxidising strength of oxyacids
Acidic strength- HClO<HClO2<HClO3<HClO4 Oxidising strength- HClO>HClO2>HClO3>HClO4
AX Type
ClF, BrF, BrCl, ICl, IBr. More polar the A-X bond greater the thermal
stability.
Linear sp3 hybridized.
AX3 type
ClF3,BrF3,ICl3,IF3 T shaped central atom sp3d hybiridzed.
AX5 type
ClF5,BrF5,IF5 Square pyramidal central atom sp3d2 hybridised
AX7type
IF7 Pentagonal pyramidal central atom sp3d3 hybridized.
Hydrolysis of Inter halogen Compounds:
Smaller halogen forms halide while the bigger halogen form oxy
halides.
ICl+H2O→HCl+HIO ICl3+2H2O→3HCl+HIO2 BrF5+3H2O→ 5HF+HBrO3 IF7+6H2O→7HF+H5IO6
CaOCl2⋅H2O
Formation :
Ca(OH)2+Cl2→Ca(OCl)Cl+H2O Pale yellow powder Oxidation state of Cl = -1 and +1
2CaOCl2⇌H2O2Ca2++2Cl−+2OCl−
Pseudohalogens
CN−,SCN−,N−3,OCN−,NCN−2 are Pseudo halide ions. Contains two or more electro-negative atoms in which one atom is
Nitrogen.
- Properties are similar to halide ions.
- Dimers are called pseudo halogens.