>SnCl - studyduniya.com Block elements - Free... · General Characteristic Valence configuration ns 2 np 2 (n= no of valence shell) Common oxidation state +2, +4 (due to inert pair

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.