Sunday, May 30, 2010


Prof. Khaqan: CHAPTER 1 BOOK 2

1- Al Razi classified elements on both physical and chemical properties. Dobriener (triads), Newland (Octaves) arranged 62 elements in ascending order of atomic masses. Mendeleev presented first regular periodic table with 8 groups and 12 periods. He placed IA & IB, IIA & IIB together.

2- Total no. of elements yet discovered = 112 s-block elements = 14 p-block =30 d-block = 40 f-block = 28. Total transition metals = 68. Total normal metals = 19. Total nonmetals =18 (gases 11 liquid 1 solids 6). Highest EA value = Cl

3- Moseley’s Periodic Law: Physical & chemical properties of elements are periodic function of their atomic numbers. In modern periodic table 8 groups & 7 periods are present. Period 1 has 2, period 2&3 have eight, period 4&5 have 18, period 6 thirty two, period 7(incomplete)has 26 elements. Elements with Z > 92 (called trans-uranic) are synthesized in laboratory. Lanthanides (period 6) & Actinides (period 7)are rare earths. Elements of group IA&IIA are alkali& alkaline earth metals, VIA chalcogens, VIIA halogens, VIIIA are noble gases.

4- Metallic character increases top to bottom in a group & decreases left to right in a period. Metals (electropositive) are good conductors of heat & electricity and form basic oxides. Nonmetals (electronegative) are bad conductors of heat & electricity and form acidic oxides. Metalloids show intermediate behavior and give amphoteric oxides.

5- Atomic sizes (atomic, ionic, covalent radii ) decrease along the period, increase down the group. Factor affecting atomic sizes are nuclear charge, number of shell & shielding effect. Positive ions are smaller than parent atom, but negative ions are larger. Greater amount of positive charge smaller the size, greater amount of negative charge bigger is the size.

6- Ionization energy & electron affinity values depend upon nuclear charge, number of shell, nature of orbital & shielding effect of inner electrons. These values decrease down the group but increase along the period. Group IIIA & VIA elements show low I.E. than IIA & VA. Group IIA, VA & VIIIA elements show abnormal E.A. values.

7- First EA value is usually negative, but higher EA values are positive. EA of second period elements is less than third period elements due to their small size.

8- Melting point /b. pt. decrease down the group for first four periods, but increase down the group in last four groups of representative elements.

9- The oxidation state of a typical element is directly or indirectly related to the group number to which the element belongs in the periodic table.

10- The electrical conductance of an element depends upon the number of free or moveable electrons.

11- Hydration energy of the ion depends upon the charge density of the ion, i.e., amount of the charge and the size of the ion. Greater is the charge density of the positive or negative ion, greater will be the heat of hydration. Value of hydration energies increases from left to right in a period, for isoelectronic species.

12- There are three types of halides: ionic, polymeric and covalent. Halides of group IA are ionic in nature, have three dimensional lattices with high melting and boiling points.

13- There are three types of hydrides formed by the elements of periodic table: ionic, intermediate and covalent.

14- Highly polar hydrides show hydrogen bonding in them

15- Oxides may be divided on the basis of their acidic, basic or amphoteric character.

16- Metallic oxides are basic in character, nonmetallic oxides are acidic in character and oxides of less electropositive elements like Zn and Pb are amphoteric (11elements).

17- Hydrogen is unique element of the periodic table. Due to similarities in properties it can be placed at the top of group IA or IVA or VIIA







Salt +

dil H2SO4

CO3-2 CO2(g)


Colorless, Odorless gas evolves with effervescence which turns lime water milky

1) Salt + water → salt soluble


i) O.S + MgSO4 sol. → white ppt. in cold

ii) O.S +CaCl2 sol. → white ppt. in cold

2) Salt+ water salt insoluble

CO3-2 confirmed (no C.T.)

HCO3- CO2(g)


Colorless, Odorless gas evolves with effervescence which turns lime water milky

Salt + water → salt soluble


i) O.S + MgSO4 sol. → white ppt. in hot

ii) O.S +CaCl2 sol. → white ppt. in hot

S-2 H2S(g)


Colorless gas evolves with rotten egg smell which turns lead acetate paper black


i)O.S+AgNO3 sol.→Black ppt

ii) O.S +Pb(CH3COO)2 sol. → Black ppt

SO3- 2 SO2(g)


Colorless gas evolves with burning “S” smell which turns K2Cr2O7 paper green (with no yellow ppt in t.t.)


i) O.S + AgNO3 sol. → white ppt

ii) O.S +BaCl2 sol. → white ppt.

iii)O.S + Pb(CH3COO)2 white ppt.

S2O3- 2 SO2(g)


Colorless gas evolves with burning “S” smell which turns K2Cr2O7 paper green (with yellow ppt in t.t.)


i) O.S + AgNO3 sol. → white ppt changing to orange brown and finally black

ii)O.S + Pb(CH3COO)2 white ppt. which blacken on heating.

NO2- NO2(g)


Reddish brown gas evolves which turns FeSO4 paper black


i) O.S + dil H2SO4 +FeSO4 sol. → Brown coloration

ii)O.S+ dil H2SO4+ KI sol + starch→ Blue coloration




Salt +

Conc. H2SO4

Cl- HCl (g)


Colorless gas evolves with pungent smell which gives fumes with NH4OH


i) O.S + AgNO3 sol. → white ppt

ii) Chromyl Chloride test

salt+ K2Cr2O7+ Conc. H2SO4(heat)

orange vapors of CrO2Cl2 .Pass the vapors in Na OH Sol.& add lead acetate sol → yellow ppts

Br - HBr (g) & Br2


Reddish brown vapors which do not turn FeSO4 paper black


i) O.S + AgNO3 sol. → pale yellow ppt


CS2 layer turns orange

I- I2


Violet vapors which turn starch paper blue


i) O.S + AgNO3 sol. → yellow ppt


CS2 layer turns purple



White fumes with pungent smell & change to reddish brown on adding paper pellets


i) O.S + Diphenylamine(few drops) → Deep blue color

ii)Ring test:

O.S +FeSO4 sol.(fresh)+Conc. H2SO4 along the side of test tube→ Dark brown ring at the junction of two solutions



Colorless vapours with vinegar smell


i) O.S + FeCl3 sol. Blood red coloration

ii)Salt + Conc. H2SO4 +C2H5OH. Pour out the product into cold water→ fruity smell of ester

iii) Palm Test: salt + oxalic acid + few drops of water on the palm & rub with finger→ vinegar like smell

(COO-)2 CO &CO2


Colorless, Odorless gas which burns on the tip of test tube &turns lime water milky

C.T. i) O.S +CaCl2 sol. → white ppt. soluble in HCl

ii) O.S+KMnO4 (acidified) → color of KMnO4 discharged on heating

Special Group

No Group Reagent

SO4-2 IDENTIFICATION: O.S + CaCl2 sol. → white ppt. insoluble in dil.HNO3

C.T. i) O.S + AgNO3 sol. → white ppt ii) O.S + Pb(CH3COO)2 white ppt

PO4-3 IDENTIFICATION: O.S + CaCl2 sol. → white ppt. soluble in dil.HNO3

C.T. i) O.S + AgNO3 sol. → yellow ppt ii) O.S + FeCl3 sol. yellow ppt.




Group Reagent

O.S + dil. HCl

Pb+2 → PbCl2 (white ppt)


White ppt. + NH4OH →not dissolved → Pb+2

C.T i) O.S + KI sol. →Bright yellow ppt.

ii)O.S+ K2CrO4 sol.→ Bright yellow ppt.

Hg2 2+ → Hg2Cl2 (white ppt)


White ppt. + NH4OH → turned black → Hg+2

C.T i) O.S + K2CrO4 sol. → Dark red ppt.

ii) O.S + KI sol. → Dirty green ppt.

Ag+ → AgCl (white ppt)


White ppt. + NH4OH → ppt dissolved → Ag+

C.T i) O.S + KI sol. → Yellow ppt.

ii) O.S + K2CrO4 sol. → Brick red ppt.

Flame Test: Make a paste of salt in conc. HCl on watch glass. Dip the clean Pt wire loop in the paste. Heat it in the oxidizing flame and note the color imparted to the flame. Bluish green = Cu2+ Crimson red = Sr2+, Green = Ba2+, Golden yellow = Na+, Brick red = Ca2+ violet = K+



O.S + dilHCl+H2S (g) → Colored ppt.

Colored ppt. + Yellow Ammonium Sulphide sol.+ Heat → ppt. insoluble

Hg +2 /Pb+2 → HgS /PbS (Black ppt.)


i) ppt.+50% HNO3 /Boil →ppt insoluble→Hg+2

ii) ppt. + 50% HNO3 / Boil → ppt soluble add dil H2SO4 → White ppt. (Pb+2)

C.T: (Same as in group I for both)

Bi+3 → Bi2S3 (Black ppt.)


Black ppt. + 50% HNO3 + Boil → ppt soluble + NH4OH in excess → White ppt. → Bi3+

C.T. i) O.S + Na OH sol. → White ppt.

ii) O.S + Na2CO3 sol. → White ppt.

Cu 2+ → Cu S (Black ppt.)


ppt. black, salt is blue → Cu 2+


i) O.S + Na OH sol. → Blue ppt.

ii) O.S + NH4OH → Light blue ppt.

Cd 2+ → CdS (Yellow ppt.)

IDENTIFICATION Yellow ppt → Cd 2+


i) O.S+ Na OH sol. →White ppt. soluble in NH4OH

ii) O.S + NH4OH → White ppt. soluble in excess



O.S + dil. HCl+ H2S (g) → Colored ppt.

Colored ppt. + Yellow Ammonium Sulphide sol. +Heat → ppt. soluble

Sn+2 → SnS (Dark brown ppt.)

IDENTIFICATION Dark brown ppt. → Sn+2

C.T. i) O.S + NaOH sol. → White ppt. soluble in excess

ii) O.S + HgCl2 sol. → White ppt. which turn black

Sn+4 → SnS (Dirty yellow ppt.)


Dirty yellow ppt. → Sn+4

C.T. i) O.S + NaOH sol. → White ppt.

ii) O.S + dil HCl + Zn metal → Grey deposit of Tin over Zn

Sb+3 → Sb2S3 (Orange yellow ppt.)


Orange yellow ppt. → Sn+3


i) O.S + NaOH sol. → White ppt.

ii) O.S + Excess of H2O → Milky sol

Borax Bead Test: (for colored salts only) Dip the loop of cleaned Pt wire in the borax, heat to make a bead. Touch the bead with salt and heat strongly. Note the color of the bead in oxidizing flame. Green in hot, blue in cold = Cu2+. Reddish brown in hot, Yellow when cold = Fe2+/Fe3+, Green in hot and cold state=Cr3+, Red in hot brown in cold = Ni2+ Blue = Co2+ & violet bead = Mn2+



O.S + NH4Cl(s) + boil + cool + NH4OH in excess

(In case of light green salt)O.S + Conc. HNO3 + boil cool NH4Cl(s) boil cool + NH4OH excess

Al+3 → Al(OH)3 Gelatinous white ppt.


Gelatinous white ppt. → Al+3


i) Lake Test:

O.S + litmus sol. + NH4OH sol. in excess → Blue ppt. floating in colorless solution

ii) O.S + Na OH sol. → White ppt.

Cr3+ → Cr(OH)3 (Dirty green ppt.)


Dirty green ppt., dark green salt → Cr3+


i) O.S + NaOH sol. → Green ppt.

ii) O.S + Na2HPO4 sol. → Pale green ppt.

Fe2+/Fe3+→ Fe(OH)3 (Reddish brown ppt.)


Reddish brown ppt. light green/yellow salt → Fe2+/Fe3+

C.T. (a) Fe2+ i) O.S + Na2CO3 sol. → White ppt.

ii) O.S + NaOH sol. → Greenish white ppt.

iii) O.S. + K3[Fe(CN)6] → Blue ppt.

(b) Fe3+ i) O.S + KCNS sol.→ Blood red color

ii) O.S + NaOH sol. → Reddish brown ppt.

Filter Ash Test: (for colorless salts only)

Shake the salt with Co(NO3)2 sol. in a test tube Dip a piece of filter paper into it. Dry and burn the filter paper. Note the color of the ash. Dirty green ash = Sn2+/Sn4+, Blue ash = Al3+/PO4 3- Green ash = Zn2+ and Pink ash = Mg2+



O.S + NH4Cl(s) + boil + cool + NH4OH in excess + H2S gas

Co2+ → CoS (Black ppt.)


ppt. Black, Dark pink salt → Co2+


i) O.S + Na OH sol. → Blue ppt.

ii) O.S. + Na2HPO4 sol. → Violet ppt.

Ni2+NiS (Black ppt.)


ppt. Black, Bright green salt → Ni2+


i) O.S + Na OH sol. → Light green ppt.

ii) O.S. + NH4OH sol. + DMG → Rose red ppt.

iii) O.S. + Na2HPO4 sol. → light green ppt.

Zn2+ → ZnS (White ppt.)


ppt. White, White salt → Zn2+


i) O.S + Na OH sol. → White ppt.

ii) O.S. + Na2CO3 sol. → White ppt.

ii) O.S. + Na2HPO4 sol. → White ppt.

Mn2+ → Mn S (Flesh colored ppt.)


Flesh colored ppt. light pink salt → Mn2+


i) O.S + Na OH sol. → White ppt which turn brown

ii) O.S. + Na2HPO4 sol. → Flesh colored ppt.



O.S + NH4Cl(s) + boil + cool + NH4OH in excess + (NH4)2CO3 sol → White ppt. Dissolve the ppt in warm CH3COOH and divide it in 3 parts.

Ba2+ → BaCO3 (White ppt.)


1st part of sol. + K2CrO4 sol. → Yellow ppt. → Ba2+


i) O.S + dil.H2SO4 sol. → White ppt.

ii) O.S. + (NH4)3PO4 sol. → White ppt. soluble in


Sr2+ → SrCO3 (White ppt.)


2nd part of sol. + Amm. sulphate sol. → White ppt. → Sr2+


i) O.S + Na2HPO4 sol. → White ppt.

ii) O.S. + Na2CO3 sol. → White ppt.

Ca2+ → CaCO3 (White ppt.)


3rd part of sol. + Amm. Oxalate sol. → White ppt. → Ca2+


i) O.S. + Na2HPO4 sol. → White ppt.

ii) O.S. + Na2CO3 sol. → White ppt.

iii) O.S. + NH4Cl+K4[Fe(CN)6] → White ppt.

Charcoal Cavity Test:

Mix a little of the salt with equal quantity of Na2CO3 and fill it in a charcoal cavity. Moist it with water and heat it strongly with a blow pipe in reducing flame. Brown encrustation = Cd2+, Reddish brown scales = Cu2+, Orange residue when hot and yellow when cold = Bi3+, Yellow residue when hot and white when cold = Zn2+



No group reagent. Each radical is identified individually


O.S. + NH4Cl(s) + boil, cool + NH4OH + Na3PO4 sol → White ppt. → Mg2+

C.T. i) O.S + NaOH sol. → White ppt

ii) O.S. + Na2CO3 sol. → White ppt.


Salt + Na OH sol heat →Ammonia Smell→ NH4+

C.T i) O.S+ Sod. Cobalti nitrite sol.→ Yellow ppt. ii) O.S + Nessler reagent → Reddish brown ppt.

iii) salt+ Na2CO3 Rub with wet finger on palm →Ammonia Smell


O.S. + CH3COOH + Sod. Cobalti nitrite sol. → No yellow ppt. → Na+

C.T i) O.S. + zinc uranyl acetate → Yellow ppt

ii) O.S + KOH sol+ Pot. Pyroantimonate→ White ppt


O.S. + Sod. Cobalti nitrite sol. → Yellow ppt. → K+

C.T. i) O.S. + tartaric acid → White ppt.

ii) O.S + picric acid → Yellow needle like crystals.


The elements of groups I-A called alkali metals and II-A called alkaline earth metals have one and two electrons in their outermost s-orbitals , so they are called s-block elements. They do not occur free in nature, due to their high reactivity. The melting and boiling points, ionization energies, electron affinities, electro negativities and heats of hydration of I-A and Il-A decrease down the group. Mg shows lowest melting and boiling points among IIA elements.

Their atomic radii, ionic radii and densities increase down the groups. Due to their low ionization energies they are good reducing agents and reducing properties increase down the group. Except Be and Mg, they give the flame tests. The oxides of I-A and II-A (except BeO) are basic in character, and give hydroxides when dissolved in water. The carbonates of I-A (except Li)are water soluble and stable towards heat. The carbonates of Li & II-A are insoluble in water and are decomposed on heating. The bicarbonates of I-A are water soluble and are decomposed on heating to give carbonates. The bicarbonates of Li & II-A exist only in the solution state.


Li differs from its family members due to its very small size, very high ionization energy and high electro negativity values from the rest of the members. Li F, Li2 CO3 , Li OH, Li2SO4 are insoluble in water. Li differs from its family members in its reactions with nitrogen, carbon and silicon, basic strength of its hydroxide, stabilities of hydroxide and the decomposition of carbonates and nitrates.

Be also differs from its family members in hardness, very high melting and boiling points and reactions with alkalies.


Oxides of I-A are more basic than those of II-A and the basic strength
of the oxides increases down the groups.

Hydroxides of these elements are alkaline in nature, but Be (OH)2 is
amphoteric and Mg(OH)2 is weakly basic.

Ca (OH)2 dissolved in water is called lime water and Mg(OH)2 dissolved in
water is called milk of magnesia.

The nitrates of Li & IIA elements decompose on heating to give metal oxides, NO2 and O2. The nitrates of IA decompose to give nitrites and O2.

The sulphates of I-A are water soluble, but the solubility of sulphates of II-A elements decreases down the group.


Sodium is prepared by Down's cell by passing the electrical current through molten NaCI and CaCl2 is added to lower the temperature to 600°C. In Down's cell the metallic fog is not produced and the material of the cell is not attacked by the products formed during electrolysis.

Caustic soda is produced from electrolysis of brine solution in Nelson's cell or diaphragm cell. Electrical current is passed through the brine solution. Chlorine and hydrogen are the by products. 11 % NaOH & 16 % NaCl is obtained in the Nelson's cell, which is concentrated to get 50% NaOH &1% NaCl .


CaSO4.2H2O is called gypsum, which is used as a fertilizer. It provides sulphur to the plants, which influences the chlorophyll development in the plant leaves. sulphur also gives improvement to the roots of the plants. Gypsum on heating gives plaster of Paris, (half-hydrated CaSO4) which is used in surgery, for making plaster walls, cast of statuary, and coins. Gypsum is added in cement to increase the setting time of the cement. When gypsum is heated at high temperature, anhydrous CaSO4 (called dead burnt gypsum) is obtained


CaO or lime is used in agriculture to neutralize the acidic soil and increase the easily-soluble phosphorus. Moreover, lime sulphur sprays are very useful and have a strong fungicidal action. Slaked lime or Ca(OH)2 is obtained, by treating lime (CaO) with water
Calcium is essential for the plant, for normal leaves development to increase the activity of micro-organisms and manage the supply of available phosphorus in the soil to the plant.

Lime is used in industry for the metallurgy of metals, and in the industries like paper, glass, sugar, bleaching powder, leather etc.

Lime is used for the white wash, preparation of acetylene and as a dehydrating reagent.

Lime mortar is prepared by mixing freshly prepared slaked lime with sand and water. It is used to bind the stones and bricks.


GENERAL CHARACTERISTICS OF ELEMENTS OF III-A Elements of group III-A are consisted of boron (B), aluminum (Al), gallium
(Ga), indium (In) and thallium (Tl). They have three electrons in the outer­
most shell. Two electrons are in the s-orbital and one unpaired electron in
the p-orbital. Melting and boiling points decrease down the group, but gallium shows abnormal behavior. Ionization energies, electron affinities and electro negativities along with the heats of sublimation decrease down the group. Boron is a semi-metal and is close to non-metallic character. Other elements of this group are metallic and good conductors of electricity.

Boron occurs in the combined state and is not an abundant element. It occurs
in traces in most of the soils. The important minerals of boron are borax or
tin cal (Na2B4O7.10H2O), colemanite (Ca2B6O11.5H2O) and orthoboric acid
(H3BO3). Boron differs from the other family members due to its extremely small size and a very high ionization energy. Boron is non-metallic in nature. B+3 ions do not exist. It mostly gives covalent compounds and its oxides are acidic in character. It has an ability to form the molecular addition compounds.

Borax can be prepared from orthoboric acid and colemanite, by treating with sodium carbonate. Borax is a white crystalline solid, its solubility increases with temperature. Above 62 C°, pentahydrated borax is obtained. Below 62 C°, decahydrated crystals are obtained.

The aqueous solution of borax is alkaline in nature and turns red litmus blue.

Borax bead test is employed by using borax to check the presence of colored
ions. They give colored beads of metal meta borates in the oxidizing and reducing flames.

Borax is used to prepare borate glass, in softening of water, in borax bead
test, in metallurgical operations, as a flux in welding, in making washing
powder, for leather industry and for the manufacture of cosmetics, soaps,
paints and medicines.

Boron gives four types of acids, which are orthoboric acid (H3BO3),
metaboric acid, (HBO2) tetra boric acids (H2B4O7) and pyroboric acids
(H6B4O9). Meta and tetra boric acids are converted to orthoboric acids when
reacted with water. Orthoboric acid can be prepared commercially from colemanite and borax. Orthoboric acid (H3BO3) is acidic in nature, but .cannot be titrated with alkalis in a usual manner. But in the presence of glycerol, titration can be done by using phenolphthalein as an indicator. It is a monobasic acid and forms the acidic solution by accepting OH- from water.

Orthoboric acid (H3BO3) is used in medicine as an antiseptic in boric
ointment and as eyewash. It finds its uses in pottery as a glaze and candle industry for stiffening of the wicks.

Aluminum is a good conductor of heat and electricity and is tarnished in the air. The reaction of aluminum with oxygen is highly exothermic and due to this
property, it is used in thermite process and photo flash bulbs.

Aluminum reacts with dilute HC1 to produce H2 gas. It reacts with concentrated H2SO4 to give SO2. Concentrated HNO3 has no
action on aluminum. Aluminum dissolves in alkalies with the evolution of hydrogen gas and soluble aluminates are produced.

Aluminum is used as an electric conductor, for making aluminum alloys, for
making utensils, frames and wires for electric transmissions. Al is also used for preventing blowholes in the metallurgy of metals.


The elements of group IV-A are carbon , silicon (Si), germanium (Ge),
tin (Sn) and lead (Pb). They have four electrons in the outermost shell. Two paired electrons are in the s-orbital and two unpaired electrons in p sub-shell.

Their melting and boiling points ionization energies, electron affinities, and electronegativities decrease down the group. However, their atomic and ionic radii increase down the group.

They show the valencies of four, but tin (Sn) and lead (Pb) can show the valency of two as well due to inert pair effect.

Metallic character increases down the group. Carbon and silicon are non-metals. Germanium is a metalloid while tin and lead are metals.

They form hydrides & halides of formula MH4 & MX4. They react with oxygen to give mono-oxides or dioxides. Silicon is a very abundant element and about 25 % of the mass of the earth’s crust is due to this element and it forms the major constituent of rocks in the form of silica and silicates. The most important compound of silicon found in the earth’s crust is SiO2 and its various forms are rock crystal, amethyst quartz, smoky quartz, rose quartz and milky quartz. Actually, it is called sand. The hydrated variety of quartz is called opal.

Carbon has three oxides as CO, CO2 and C3O2. In CO, two bonds are covalent and one is a coordinate covalent bond. CO is a polar molecule and its dipole moment is 0.112 D. CO acts as a good ligand.

CO2 has a linear structure in which carbon is sp-hybridized. The net dipole moment of CO2 is zero and the solid CO2 called dry ice, has a face centered cubic structure. CO2 is acidic in nature.

SiO2 is transparent to light, and can tolerate high temperature. It has low thermal expansion. It is an excellent insulator. It is hard, brittle and elastic.

SiO2 is insoluble in water. It is not affected by acids except HF.

Silicates are the derivatives of meta-silicic acid (H2SiO3). Na2SiO3 is also called water glass, or soluble glass. It is soluble in water. Its aqueous solution is strongly alkaline. It is used for the preparation of chemical garden.

Sodium silicate is used as filler in soap manufacture, in textile as a fire
proof, for furniture polish, in calico printing and as a preservative of eggs.

Aluminum silicates are used to make porcelain and china wares and to glaze the stone wares.

Talc or soap stone is another important silicate with the formula
Mg3H2(SiO3)4. It is greasy to touch and is used in making cosmetics and
some Household articles.

Asbestos is a mixed silicate of calcium and magnesium with the formula
CaMg3(SiO3)4. It is used to make incombustible fabrics and hard board etc.

The lower silicones are oily liquids while higher members are waxy solids.
The viscosity of silicon oils doesn’t change that much with temperature as compared to the petroleum oils.

Methyl silicones of high molar masses resemble rubber and are used as
insulating materials for electrical motors. Silicones are used to cover the
surfaces of plastics, asbestos, glass, leather, filter paper and blotting paper.

Germanium, selenium and silicon are important semiconductors. Electrical
conductivity of semi-conductors increases by increasing temperature.

Semi-conductors are used in transistors, which are used in radio, T.V.
computers and calculators.


Pb2O is a black powder. PbO, which is called litharge or massicot and exists
in two crystalline forms i.e. rhombic, (yellow) and the tetragonal (red). When litharge is boiled with water and olive oil, we get lead oleate, a sticky adhesive mass. Litharge is used in preparing flint glass and paints, for preparing oils and varnishes, for the manufacture of glazes, in rubber industry and lead storage battery.

Pb3O4 called the red lead, minium or sandhur, is scarlet in color.Pb3O4 is used as a red pigment in paints, in glass industry, in lead storage batteries

PbO2 is a reddish brown powder, less soluble in water and dissolves in alkaline water to give soluble plumbates.

The basic lead carbonate 2PbCO3Pb (OH)2 is a white amorphous powder insoluble in water. When white lead is mixed with linseed oil, it has a good covering power. Lead chromate is used as a pigment under the name of chrome yellow. When it is mixed with lead sulphate or barium sulphate, it is used as a yellow pigment.


The elements of group VA are nitrogen, phosphorus, arsenic, antimony and

bismuth. Except nitrogen, all elements of group V-A occur in the combined state. Nitrogen constitutes approximately 80 % of the earth's atmosphere. All the elements of group VA show allotropy except bismuth. Nitrogen is a gas while other substances are solids. The melting and boiling points steadily increase, but Bi shows abnormal behavior. The ionization energies, electron affinities and electro negativities decrease down the group. Nitrogen and phosphorus are non-metals. As and Sb are metalloids while Bi is a metal. The elements of group V-A give the oxides of type M2O3, M2O4 and M2O5 but nitrogen also gives N2O and NO.The elements of group V-A give tri halides and penta halides of general formula MX3 and MX5. However nitrogen cannot give penta halides.

NITROGEN differs from its family members in many respects. It exists in free state In diatomic form and is chemically inert as compared to others.

Nitrogen gives five oxides i.e. N2O, NO, N2O3, NO2 and N2O5.

N2O can be prepared by reaction of Zn and dil. HNO3 or by heating ammonium nitrate.N2O is a colorless gas with a sweetish taste. It is also called laughing gas. It
is heavier than air and neutral to litmus.
N2O is a supporter of combustion and supplies, oxygen to P, Mg, Na, Cu, hydrogen and NH3. In all these reactions, N2 gas is set free.

NO gas is prepared by treating Cu with cold dilute HNO3.

NO is a colorless gas, heavier than air, sparingly soluble in H2O and neutral to litmus.

NO acts as an oxidizing agent and a reducing agent as well. It is used as a catalyst in lead chamber process for manufacture of H2SO4.

NO2 gas can be prepared by the reaction of Cu metal with cone. HNO3 or by
heating Pb(NO3)2. NO2 is a reddish brown gas, soluble in waterand forms blue acidic solution. NO2 reacts with non-metals and converts them into their oxides. It gives acid
by reacting with water and combines with alkalies to give salt and water.Moreover, it is a good oxidizing agent. NO2 is used as a rocket fuel and a starting material in manufacture of HNO3. HNO2 can be prepared by the reaction of N2O3 with water and by treating
barium nitrite with dil. H2SO4.

HNO2 is pale blue in color due to the presence of N2O3. It is a weak acid and is an unstable compound.

HNO2 decomposes in cold solution and undergoes auto oxidation. It acts as
an oxidizing agent as well as a reducing agent. It also combines with urea and aniline to give nitrogen gas.

HNO3 can be prepared on industrial scale by Birkland Eyde process in which nitrogen and oxygen are converted to NO by electrical current at 3000°C. NO is converted to NO2 at 600°C which reacts with H2O to give HNO3.

HNO3 is an oxidizing substance and oxidizes non-metals like C, S, P, I and
Si. It also oxidizes metalloids like As and Sb.

HNO3 reacts with metals. The nature of products depends upon the
concentration of acid, temperature of reaction and nature of metal.
it oxidizes FeSO4, H2S, HI

A mixture of concentrated HC1 and concentrated HNO3 in the ratio of 3 : 1
is called aqua regia and can dissolve the noble metals like Au and Pt as
their chlorides.

HNO3 is used in the manufacture of explosives, preparation of artificial silk,
purification of Au and Ag and as an important reagent in laboratory.

Red phosphorus is obtained by heating white phosphorus. It exists in the'
form of macro molecules. Black phosphorus is obtained by heating red
phosphorus to high temperature and pressure.

Phosphorus gives two types i.e. PX3 and PX5. PC13 is obtained by
reaction of-phosphorus with C12 or by treating phosphorus with SOC12.

The structure of PCI3 is pyramidal just like NH3, but the bond angles are just
less than those of NH3.

PC13 is a colourless liquid but PC15 is a yellowish white crystalline solid.
The structure of PC15 is trigonal bipyramid and P completes its ten electrons
in its outermost valence shells.

Phosphorus gives variety of oxides. P2O3 is prepared by the reaction of
phosphorus with limited supply of oxygen or by reacting phosphorus with

P2O3 combines with H2O to give H3PO3. It is a white wax-like solid and has
a garlic smell. It is a highly poisonous substance.

P2O5 can be prepared by the reaction of phosphorus with oxygen or by
reacting phosphorus with CO2.

P2O5 is a white powdery solid. It gives strong phosphorescence after
illumination. It is a volatile solid and sublimes at 360°C.

It reacts with water to give H3PO4 and is a very strong dehydrating reagent.
It is extensively used as a drying and dehydrating reagent in the laboratory.

Phosphorus gives various oxyacids, but H3PO3 and H3PO4 are important.
H3PO3 is prepared by the reaction of P2O3 or PC13 with water.

H3PO3 is a white solid and is highly soluble in water. It hasv reducing
properties and can reduce CuSO4, AgNO3 and KMnO4.

H3PO4 can be prepared by the reaction of phosphorus pentaoxide with water.

H3PO4 is a hard substance having rhombic prisms and soluble in water in all

H3PO4 is a tribasic acid and is influenced by the temperature.

H3PO4 is used as a nerve tonic, in the manufacture of fertilizers and for the
manufacture of dyes and enamels.



Elements of group VI-A are O, S, Se, Te and Po. They have six electrons in
the outermost principal quantum number, Two electrons are in s-orbital and
four in p-orbitals.

They have 2 unpaired electrons in p-orbitals and can show the oxidation
states of +2. But if they do the unpairing of electrons, then the oxidation
numbers can become +4 and +6. This is due to availability of d-orbitals.

Oxygen cap not do tTie unpairing of electrons and mostly shows the valency
of two.

Their melting and boiling points increase down the group, T^ut phosphorus shows abnormal behaviour. Their densities also increase down the group.

Their ionization energies and electron affinities decrease down the group.

Oxygen exists in the form of O2 molecule. Other elements of this group are octa-atomic and have cyclic structures.

Almost all the elements of this group show allotropy.

Oxygen shows the abnormal behavior in its family members. Oxygen is the most widely distributed and is common among all elements. It is 46.6 % of our earth's crust.

Sulphur occurs in the form of ores and minerals and also occurs in many organic compounds in animal and vegetable products like onions, garlic, mustard, hair, oils, eggs and proteins.

H2SO4 is prepared on industrial scale by lead chamber process and contact process. In contact process, the SO2 gas is purified to the maximum extentand a solid catalyst like V2O5 or Pi. metal is used. The temperature is maintained utp 400 - 500°C. SO3 gas so produced is dissolved in H2SO4 to get oleum, which is diluted to get H2SO4.

Pure H2SO4 is a colourless, oily liquid without odour. When dissolved in
water, it liberates large amount of heat. It is highly corrosive to skin.

H2SO4 acts as a dehydrating agent and acts as a very strong oxidizing agent.
It reacts with KMnO4 and K2Cr2O7 and releases atomic oxygen. So, this
mixture can be used as an oxidizing agent.

H2SO4 is the king of chemicals. It is a barometer of the industries of a country. It is used in manufacture of fertilizers, refining of petroleum,
manufacture of HC1, HNOs and H3PO4, dyes, drugs, plastics, disinfectants, paints and synthetic fiber.

Moreover, H2SO4 is used as a dehydrating agent, in lead storage batteries and as an important laboratory reagent.

Chapter 5 Halogens & Noble Gases

Fluorine shows the peculiar behavior in its family members due to its very small size, high electro negativity, very low bond dissociation energy and non-availability of the d-orbital for the promotion of electron.

Halogens are oxidizing substances and the factors like heat of dissociation, electron affinities, hydration energies and heat of vaporization control the oxidizing properties. Due to great oxidizing capabilities of fluorine and chlorine, they can act as decolorizing agents and can decolorize Litmus and universal indicator. Chlorine can be used in bleaching powder.

HF & HCl can be prepared by the reaction of their salts with concentrated sulphuric acid HBr & HI not be prepared by this method. HF & HI can not be prepared directly by the combination of H2 and X2. The melting and boiling points heats of fusion and heats of vaporization of halogen acids gradually increase from HCl to HI, but HF shows exceptional behavior due to hydrogen bonding. Their bond dissociation energies decrease from HF to HI.

HF is a colorless volatile liquid, highly soluble in water, bad conductor of electricity and has strong hydrogen bonding. HF attacks on glass and it is handled in TEFLON. When HF is absolutely dry, it is stored in copper or stainless steel vessels. In vapor state HF molecules consist of an equilibrium mixture of monomers of HF and its hexamer (HF)6. There are chains and rings of HF molecules.

The acid strengths and reducing properties of Halogen acids increase from HF to HI Oxides of chlorine are generally unstable. C1O2 is prepared by the reduction of NaClO3 with NaCl or by treating KC1O3 with sulphuric acid and oxalic acid.

C1O2 is a dark yellow pungent smelling gas and paramagnetic in nature.

C1O2 is used as an antiseptic, for the purification of water, to bleach the
cellulose material, to improve the quality of low grade fats and oils and to
make the wool unshrinkable.

C12O7 is slowly reacts with water giving HC1O4 and is called anhydride of perchloric acid. When chlorine reacts with NaOH at 15°C, NaCl and NaCIO are produced. When the reaction is carried out at 70°C, NaCl and NaClO3 are produced. In these reactions, called disproportionation reaction the chlorine atom in C12 is oxidized and reduced simultaneously. Halogens give oxy acids of formulas HXO, HXO2, HXO3 and HXO4. . Fluorine also does not give such oxy acids.

The oxy acids are unstable compounds and they cannot be isolated in the pure form. They are stable in aqueous solutions or in the form of their salts. The hydrogen atoms in oxy acids are connected with halogens through oxygen atoms. Greater the number of oxygen atoms greater the acid strength and the thermal stability and weaker is oxidizing agent. The oxy acids of chlorine are stronger than those of bromine which are stronger than that of iodine with the same number of oxygen atoms.

HCIO4 is prepared by the reaction of KClO4 with concentrated H2SO4. Bleaching powder Ca Cl (O Cl) is prepared by the reaction of slaked lime with C12 in Hasenclever method or Beckmann's method. Bleaching powder is always packed in airtight containers in order to avoid the lose of chlorine. Bleaching powder is a yellowish, white powder and has a strong smell of
chlorine. It acts as an oxidizing agent.due to the presence of OC1- ion in water. Bleaching powder reacts with dilute and concentrated acids to give various types of products. When' it reacts with ammonia, it releases nitrogen gas with the formation of CaCI2.The value of the available chlorine with a good sample of bleaching powder is in the range of 35 —40 %.

Bleaching powder is used for the manufacture of chloroform on industrial scale, as a disinfectant, for making un-shrinkable wool, for bleaching cotton, paper, pulp etc. Chloro-fluorocarbons, Freons are non-toxic and non-inflammable and are used in refrigeration and air conditioning .

Terta-fluoroethylene can be polymerized to give Teflon. It is a plastic type inert material, is not soluble in any solvent and not attacked by strong acids,
alkalies and oxidizing agents.

Teflon, is used in the construction of chemical plants and due to high
electrical resistance it is used as an insulating material in cables. Halothane
is used as an anesthetic.

Chlorine is used in the manufacture of bleaching powder and as a
disinfectant in swimming pools. It is used for the manufacture of antiseptics,
insecticides, weed killers and herbicides. Moreover, chlorine is used for the
preparation of HCI, PVC, CHCl3 and CCl4.

Bromine is used for the preparation of ethylene dibromide, which is used in
petrol. It acts as a good germicide and fungicide. AgBr is used in

Iodine is used in pharmaceutical industries as a disinfectant and germicide.
Tincture of iodine is also prepared. Iodine is analgesic and is used for the
treatment of thyroid. Elements of group VIII A are helium (He), neon (Ne), argon (Ar), krypton
(Kr), xenon (Xe) and radon (Rn). They are called noble gases or inert gases.

Their outer most shell is complete having eight electrons except helium,
which has only two electrons. They are least reactive due to complete octet.

Helium is the second most abundant element in the universe after hydrogen. All the noble gases are colorless and odorless, but when subjected to high voltage, they give different colors in the spectra. Neon glows reddish and its discharge is most intense at ordinary voltage and current. Krypton has brilliant green and orange spectral lines.

Their melting and boiling points increase down the group due to the
increasing polarizabilities. Their ionization energies decrease down .the
group. They are least soluble in water but solubility increases down the
group. Their heats of vaporization increase down the group. Helium is used in the treatment of asthma, filling of weather balloons,
producing low temperature, preservation of food, producing inert
atmosphere, as a shielding gas in signal lights and in atomic reactors.

Neon is used in making neon-advertising signs, in high voltage indicators
and T.V. tubes and in making glass lasers.

Argon is used for filling the electrical bulbs, in Geiger counters, for
producing inert atmosphere, filling of fluorescent tubes and radio valves. By
mixing with neon in neon signs, the light of varying colors can be obtained.

Krypton is used for filling fluorescent tubes in flash lamps and for the
measurement of thickness of sheets.

Xenon is used in bactericidal lamps and in atomic energy field in bubble


(1) d and f-block elements are called outer transition and inner transition elements respectively. They are all called transition elements, because their properties are in between s and p-block elements. The general electronic configuration of outer transition elements is:(n-1)d1-10 ns1-2 and they contain two types of valence electrons—the d electrons and the s electrons.

(2) The electronic distribution of Cr and Cu family has one electron in their s-orbitals and 5 and 10 electrons in d-orbitals respectively. The elements of II-B (Zn, Cd, Hg) and the elements of III-B (Sc, Y, La) are called non-typical transition elements.

(3) Transition elements have similar physical and chemical properties. For example, they are all hard and strong metals, play important role in industries, form alloys, show variable oxidation states, and form colored compounds and paramagnetic ions.

(4) Their melting and boiling points are closely associated with their binding energies. The melting and boiling points along with the binding energies increase from left to the right up to the middle of the series and then decrease to minimum towards the end of the series. This is all due to the varying number of-unpaired electrons in their valence shell.

(5) Transition metal ions are mostly paramagnetic in nature due to the presence of unpaired electrons. Their magnetic moment is measured in unit called Bohr’s magnetrons. It increases from left to the right up to the middle of the series and then decreases to minimum towards the end of the series The ions with highest paramagnetic property are Mn+2,Fe+3. Transition elements also show variable oxidation states due to the involvement of d electrons. This property increases up to the middle of the series and then decreases onwards. The highest o.s. is 7 which is associated with Mn family(VII B group)

(6) The compounds of the transition elements are colored, which is associated with the presence of unpaired electrons in the d-orbitals, which have lost their degeneracy.

(7) Transition elements can form interstitial compounds by accommodating small sized atoms like hydrogen, oxygen, carbon, boron etc. in their interstices. These are non-stoichiometric (not true) compounds. They also give valuable substitutional alloys with other transition elements.

(8) Transition elements make the complex compounds by acceptance of certain number of electron pairs from donor species, called ligands. Transition metals or ion is surrounded by negative or neutral ligands
to form a complex ion. This complex, written within the brackets called the co-ordination sphere, may have positive or a negative charge or even neutral. The number electron pairs donated to the central transition metal or ion is known as the co-ordination number.

(9) Ligands can be monodentate, bidentate or polydentate. The charge on the co-ordination sphere is the sum of the charge on the central metal atom or ion and charge of the ligands. Chelates are complexes of polydentate ligands.

(10) While giving IUPAC names to the complex compounds, cation are written before the anions. Names of negatively charged ligands end in o,while the neutral ligands are called as such, for example, Cl-,NO2-,CN-,OH-,CO3-,C2O4-, are called chloro, nitro, cyano, hydroxo, carbonato, and oxalato. The neutral ligands NH3,H2O, CO are called ammine, aqua and carbonyl respectively.

(11) The ligands are named alphabetically with prefixes as di-for two, tri-for
three, tetra-for four, penta-for five and hexa for six etc. If the complex ion is negatively charged, the name of the metal ends in -ate. The oxidation state of the metal ion is represented by the Roman numeral in square bracket.

(12) The geometry of the complex ion depends upon the nature of the hybridization of central metal atom or ion. For sp3-hybridization, the complex ion is tetrahedral. For dsp2, it is square planar, for dsp3, trigonal bipyramid and for d2sp3, it is octahedral.

(13) The important ores of iron are Limonite (Fe2O3.3H2O), hematite (Fe2O3) and magnetite (Fe3O4). The three commercial forms of iron are: cast iron or pig iron (2.5 - 4.5 % carbon) wrought iron (0.12 - 0.25 % carbon) and steel (0.25 - 2.5 % carbon).

(14) Cast iron is converted to wrought iron, by using puddling furnace. In this
furnace, the iron oxide of hematite lining is reduced to iron metal and the impurities like carbon, sulphur are oxidized to volatile oxides, while manganese, phosphorous and silicon are converted to their silicates and phosphates and separated as slag.

(15) Steel can be classified as: mild steel (0.1%-0.2%C) , medium steel (0.2%-0.7%C)and high carbon steel(0.7%-1.5%C).

(16) Steel can be manufactured by Open hearth process or Bessemer's process. Open hearth furnace works on the regenerative principle of heat economy. The furnace can have acidic lining of SiO2 or basic lining of dolomite (CaO. MgO). The impurities present in the pig iron and scrap steel are oxidized to CO,SiO2, MnO, SO2 and P2O5, which react with different metal oxides to form slag.Slag also contains Ca3(PO4)2. Slag, which floats on the surface of molten metal is removed and ferromanganese (Fe, M n, C) is added to get the steel of the required quality. The whole process takes 10 hours.

(17) In Bessemer process, pig iron or cast iron is directly taken from the blast
furnace and the blast of hot air is injected through the holes at the bottom. Carbon, silicon and manganese are converted to their oxides and finally to silicates. Carbon monoxide produced burns at the mouth of the converter with a blue flame. When the blue flame subsides, a mixture of carbon, iron and manganese is added to get the steel of the required quality. In order to remove entrapped bubbles of gases, a little aluminum is added. The addition of manganese desulphurizes the steel and increase its hardness and tensile strength.

(18) Corrosion is the chemical decay of metal due to the action of surrounding. During the process of corrosion by air the surface of the metals is coated with compact layer of oxides, sulphides and carbonates. This layer protects the metal from further attack. In the presence of water, the corrosion, however, penetrates into the metal. The electrochemical theory successfully explains the phenomenon of corrosion. Coating of other metals, alloying, coating with paint and varnish can prevent corrosion.

Coating of less reactive metals is called cathode coating, e.g., tin plating. If the protective coating of tin is damaged, iron comes in contact with the moisture and corrosion becomes more rapid. Galvanizing or anode coating is the coating of more reactive metal, e.g., zinc coating.

(19) K2CrO4 is prepared by the oxidation of KCrO2 with bromine and KOH. or Cr2O3 with KClO3 in basic medium. It can also be prepared by heating natural chromite with K2CO3 and oxygen. When K2CrO4 is treated with sulphuric acid, equilibrium is established between chromate and dichromate ions. PbCrO4 (lead chromate) is a yellow pigment, used to prepare yellow crown. The structure of CrO4-2 is perfectly tetrahedral.

(20) K2Cr 2O7, a water soluble orange red crystalline solid, is prepared by the reaction of K2CrO4 with Sulphuric acid, or by the reaction of Na2Cr 2O7 with KCl.K2Cr2O7 is a very powerful oxidizing agent in the presence of Sulphuric acid. Acidified K2Cr2O7 can oxidize FeSO4, KI, FeSO4 and H2S.Acidified K2Cr2O7 on treatment with chloride ion, produces a deep red liquid called chromyl chloride. This is one of the tests for chloride radical. K2Cr2O7 is used in the dyeing, as an oxidizing agent and in leather tanning.

(21) KMnO4 is prepared by the reaction of K2MnO4 with sulphuric acid or with chlorine (stadeler’s process), with CO2 or by its electrolytic oxidation. On large scale K2MnO4 is prepared from the oxidation of pyrolusite(MnO2). KMnO4 has dark purple lustrous crystals and gives deep pink color in the solution state, with low solubility in water. KMnO4 is a powerful oxidizing agent in the presence of H2SO4 and can oxidize H2S, FeSO4, and oxalic acid. KMnO4 is used as an oxidizing agent, disinfectant and germicide and in the manufacture of organic compounds.The structure of KMnO4 is perfectly tetrahedral.

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