p block element

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P-BLOCK ELEMENTS

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P BLOCK ELEMENT

The elements in which last electron enters in p orbital are known as p block elements. As in a p sub-shell there are three orbital which can accommodate six electrons, so p block elements includes six columns i.e. 13 18 groups. The electronic configuration of p block elements in valence shell is ns2np1-6. The inner core of electronic configuration may differ. The difference in inner core greatly influence their physical and chemical properties.

Oxidation state Maximum oxidation state shown by p block elements is equal to the number of valence electrons. So number of possible oxidation states increase from left to right in the period. This is known as group oxidation state. Some

element also shows oxidation state two unit less than group oxidation state. This tendency increases from top to

bottom in a group and is known as inert pair effect.

Metallic & Non-metallic Character

p block elements include metals, non-metals and metalloids. Metallic nature increases from top to bottom in each group and decreases from left to right in each period. Non-metallic nature increases from left to right in each period

and decreases from top to bottom in each group. The compounds formed by highly reactive metals with highly

reactive non-metals are ionic because of large difference in their electro negativities. Compounds formed between

non-metals are largely covalent due to small difference in their electro negativities. The oxides of metals are basic in

nature where as oxides of nonmetals are acidic in nature.

The first member of the group differs from other members of the group due to their 1. Abnormally small size 2. High electro negativity 3. Absence of nd sub-shell.

2nd period elements of P block (1st member). Shows maximum covalency of 4. Elements of third period has 3d sub-shell. So valence shell may be expanded i.e. electrons from 3s and 3p sub-shell may be excited to 3d sub-shell.

So they show higher oxidation states also. 2nd period elements due to their small size, have a tendency to form

pp bonds (multiple bonds) C C, C C, N N.

The heavier elements also form bonds (d p or d d) as d orbitals have higher energy so these bonds contribute less to the overall stability.

GROUP 13 ELEMENTS (THE BORON FAMILY) This group includes Boron, Aluminum, Gallium, Indium and Thallium. Boron is non-metal, aluminium; Gallium,

Indium and Thallium are metals. Occurrence: Boron Orthoboric acid (H3BO3), Borax (Na2B4O7.10H2O), Kernite (Na2B4O7.4H2O). Boron has two isotopes B10 (19%), B11 (81%) Aluminum is most abundant metal in earth crust. Main ores are

(1) Bauxite Al2O3.2H2O (2) Cryolite Na3AlF6. Gallium, Indium and Thallium are less abundant.

1. Electronic configuration: They have ns3 np1 electronic configuration in outermost shell. (a) Boron (B) 5 [He] 2s2, 2p1 (b) Aluminum (Al) 13 [Ne] 3s2, 3p1 (c) Gallium (Ga) 3 [Ar] 3d10, 4s2, 4p1 (d) Indium (In) 49 [Kr] 4d10, 5s2, 5p1 (e) Thallium (Tl) 81 [Xe] 4f14, 5d10, 6s2, 6p1 (f)

2. Atomic radii: Atomic radii / Ionic radii are smaller than those of corresponding earth metals. Atomic radii / Ionic radii increases from top to bottom due to addition of one shell in each successive period. Atomic radius of

gallium is less than that of Aluminum due to poor shielding effect of d electrons.



3. Ionization Enthalpy: 1st I.E. of 13th column elements is lower than 1st I.E. of alkaline earth metal of same period. Decrease I.E. from B to Al is due to increase in size. Increase in 1st I.E. between Al & Ga and between In &

Tl due to low screening effect of d & f electrons. The order of ionization enthalpies are iH3 > iH2 > iH1

4. Electro negativities: Electro negativity decreases from B to Al and then increases marginally.

PHYSICAL PROPERTIES

Boron is non-metallic in nature, extremely hard and black coloured solid. It has a strong crystalline lattice with very

high melting point. Rest of the members are soft metals with low melting point and high electrical conductance.

Gallium has low melting point (303 K) but very high boiling point (2676 K). Density increases from Boron to

Gallium.

1. Oxidation State Due to abnormally small size of Boron, the sum of its first three ionization enthalpies is very high, so boron does not

form ionic compounds but it forms covalent compounds, in which it shows an oxidation state of III. In aluminum

sum of first three I.E. is low, so Al forms Al+3 ions and aluminum is a strong electropositive element.

In higher members of the group due to poor shielding effect of d and f electrons, the increased effective nuclear

charge holds ns electrons tightly (inert pair effect). So only p orbitals may be involved in bonding showing an O.S. of +1. Ga, In and Tl show both +1 and +3 O.S. Tendency to show +1 O.S. increases from Ga to Tl. Compounds

in +1 O.S. are more ionic than in +3 O.S. In trivalent state the number of electrons around the central atom will be

only six, so these compounds are electron deficient so act as Lewis acid tendency to behave as Lewis acid decreases

from top to bottom. BCl3 easily accepts electrons from NH3 to form BCl3.NH3.

AlCl3 achieves stability by forming a dimmer.

In trivalent state most of the compounds are covalent and are easily hydrolysed to give tetrahedral [M(OH)4]

with sp3 hybridisation. AlCl3 exists as dimer in organic solvents but in aqueous solution behaves as ionic compound with

the formation of octahedral [Al(H2O)6]3+ ion in which Al is sp3d2 hybridised.

2. Reactivity with air Boron is unreactive in crystalline form, amorphous boron on heating in air gives B2O3. Aluminum forms a thin layer

of oxide on the surface which protects the metal from further attack. On heating in air Al forms Al2O3 4M + 3O2 2M2O3 When heated with nitrogen at high temperature they form nitrides.

3M + N2 M3N2.

3. Nature of Oxides Oxide of boron is acidic in nature, aluminum and gallium oxides are amphoteric and indium and thallium oxides are basic in nature.

Semi precious stones like ruby, sapphire, amethyst are Al2O3 and different colours are due to inpurities

4. Reaction with Acids Boron does not react with acids. Aluminum dissolves in mineral acids to give salts and hydrogen.

2Al (s) + 6HCl (aq) 2Al3+ (aq) + 6Cl (aq) + 3H2 (s) In conc. HNO3 aluminum becomes passive due to formation of a protective oxide layer on the surface.



5. Reaction with alkalies Boron does react with aqueous alkali. Aluminum reacts with aqueous alkali to liberate hydrogen gas.

2Al (s) + 2NaOH (aq) + 6H2O (l) 2Na+[Al(OH)4]

(aq) + 3H2 (g)

(Sodiumtetrahydroxoaluminate)

6. Reaction with halogens React with halogens to give trihalides.

2M (s) + 3X2 (g) 2MX3

ANOMALOUS BEHAVIOUR OF BORON

As Boron has very small atomic size and high ionization enthalpy, so it shows properties different from other

members of the group.

1. B is a non-metal other members of the group are metals. 2. B shows maximum covalency of four due to absence of nd sub-shell rest of the members show maximum

covalency of six.

3. Only boron shows allotropy. 4. Boron does not react to from cation in aqueous solution as the sum of first three ionization enthalpy has

very high value. 5. Boron halides are monomeric where as halides of other members of the group are dimeric. 6. Boron does not react with water or steam where as other members of the group react. 7. Oxide of Boron is acidic in nature. 8. Boron is not attacked by non oxidizing acids. 9. Boron dissolves in conc. HNO3 to give H3BO3, other members of the group become passive.

SOME IMPORTANT COMPOUNDS OF BORON

BORAX: Na2[B4O5(OH)4.8H2O] (Na2B4O7.10H2O) Dissolves in water to give an alkaline solution as it is a salt of weaker acid and strong base.

Na2B4O7 + 7H2O 2NaOH + 4H3BO3

Effect of heat on borax On heating borax first loses water molecules and swells up. On further heating it turns into transparent liquid which

solidifies into sodiummetaborate and then boric anhydride.

Na2B4O7.10H2O

Na2B4O7

2NaBO2 + B2O3 (Boric Anhydride)

(sodiummetaborate)

Metaborate of many transition elements are coloured Co(BO2)2 is blue in colour. (blue bead)

ORTHOBORIC ACID H3BO3

Preparation: By acidifying an aqueous solution of borax. Na2B4O7 + 2HCl + 5H2O 2NaCl + 4B(OH)3 [Orthoboric acid] Orthoboric acid is a white crystalline solid with a soapy touch. Sparingly soluble in cold water but highly soluble in

hot water. Boric acid is a weak monobasic acid. It is not a protonic acid, it is a weak Lewis acid as it accepts

electrons from hydroxyl ion.

B(OH)3 + 2H2O [B(OH)4] + H3O

+

Effect of Heat

On heating othoboric acid at about 370 K it gives metaboric acid which on further heating gives B2O3.

H3BO3 OH 2

HBO2

B2O3

[Metaboric acid]

Orthoboric acid has a layer structure in which planar BO3 units are joined by H bonds.



Structure of Boric acid

DIBORANE B2H2

Preparation 1. By the action of LiAlH4 with BF3 in diethyl ether

4BF3 + 3LiAlH4 2B2H6 + 3LiF + AlF3 2. By the oxidation of sodium borohydride with iodine.

2NaBH4 + I2 B2H6 + 2NaI + H2 3. By the reaction of BF3 with sodium hydride.

2BF3 + 6NaH K 400

B2H6 + 6NaF

Properties: It is a colourless highly toxic gas, Boiling point 180 K. Diborane catches fire spontaneously when

exposed to air.

1. Burns in air to give B2O3 with release of enormous amount of heat.

B2H6 + 3O2 B2O3 + 3H2O, c Ho = 1976 kJ mol1.

2. Reacts with water to give boric acid.

B2H6 (g) + 6H2O (l) 2B(OH)3 (aq) + 6H2 (g) 3. It acts as a Lewis acid and combines with Lewis bases.

B2H6 + 2NMe3 2BH3.NMe3 B2H6 + 2CO 2BH3.CO 4. Reaction with Ammonia

B2H6 + 2NH3 [BH2(NH3)2]+ [BH4]

On heating with ammonia diborane gives borazine.

(B3N3H6) known as inorganic benzene

3B2H6 + 6NH3 3[BH2(NH3)2]+ [BH4]

Heat

2B3N3H6 + 12H2

Structure of borazine is

Structure of diborane

Two boron atoms and four hydrogen atoms lie in one plane. There are two bridged hydrogen atoms, one above the

plane and one below the plane. Four BH bond which are in same plane are regular two centre, two electron bonds while the bridged BHB bonds are three centre, two electron bonds.



Orbital structure of diborane

Both boron atoms are sp3 hybridised. Three hybrid orbitals have one electron each and fourth is empty. With the

help of two hybrid orbitals each boron atom forms normal B H bonds. With the help of one half filled atomic orbital of one boron atom and vacant atomic orbital of 2nd B atom and 1S atomic orbital of Hydrogen, overlap to give 3 centre 2 electron bond also known as Banana bonds.

BOROHYDRIDES

When diborane in diethyl ether is treated with metal hydrides, borohydrides are obtained.

2MH + B2H6 2M+[BH4]

[ = Li or Na]

LiBH4 and NaBH4 are used as reducing agents in organic synthesis.

Uses of Boron & Aluminium

1. Boron fibres are used in making bullet-proof vest and light composite material for aircraft. 2. B10 has high ability to absorb neutrons, so metal borides are used as protective shield and control rods in

nuclear reactors.

3. Borax and boric acid are used in making heat resistant glasses, Glass wool and fibre glass. 4. Borax is used as a flux. 5. Boric acid is used as a mild antiseptic. 6. Aluminium is a bright silvery white metal with high tensile strength. It has high electrical and thermal

conductance. It forms alloys with Cu, Mn, Mg, Si and Zinc.

7. Alloys are used in making utensils, construction, aeroplane etc.

BORON TRIHALIDES

Boron reacts with halogen to give trihalides. BX3 (X = F, Cl, Br, I) which act as a Lewis acid. In trihalides B is in sp2 hybrid state and is planar. Vacant Pz A.O. of Boron overlaps with Pz A.O. of halogen having a lone pair of

electrons (back bonding), so the BX bond is shorter than expected. This also decreases Lewis acid character of Boron trihalides. With increase in size of halogen atom, the back bonding decreases so acidic character increases. So

Lewis acid character of Boron halide is

BI3 > BBr3 > BCl3 > BF3



14

TH COLUMN ELEMENTS CARBON FAMILY

14th column includes Carbon (C), Silicon (S), Germanium (Ge), Tin (Sn) and lead (Pb). Carbon is distributed in

nature in free as well as in combined state. In elemental state carbon is present as Coal, Graphite, Diamond. In

combined state carbon is present as metal carbonates, hydrocarbons and carbon dioxides. In combination with other elements carbon is a component of living tissues, drugs etc. there are two stable isotopes of carbon C12 and C13 and

radioactive C14.

Silicon is 2nd most abundant element in the earth crust in the form of silica and silicates. Silicate is important

component of ceramics, glass and cement. Germanium exists only in traces. Tin occurs as tinstone or Casseterite

(SnO2) lead as galena (PbS). Ultrapure form of Germanium and Silicon are used in making transistors and

Semiconductors.

Properties:

1. Electronic configuration Carbon (C) (6) [He] 2s2, 2p2

Silicon (Si) (14) [Ne] 3s2, 3p2

Germanium (Ge) (32) [Ar] 3d10, 4s2, 4p2 Tin (Sn) (50) [Kr] 4d10, 5s2, 5p2

Lead (Pb) (82) [Xe] 4f14, 5d10, 6s2, 6p2

2. Covalent Radii Atomic radii of 14th column of elements are smaller than those of 13th column of respective period due to

increase in nuclear charge. Atomic radii increases from C to Pb due to addition of one shell in each period.

There is considerable increase from C to Si, then there is small increase from Si to Pb due to poor Shielding

effect of d and f electrons.

3. Ionization Enthalpy The 1st ionization enthalpy of group 14 elements are higher than those of group 13 of the same period due to

increase in nuclear charge. 1st ionization enthalpy decreases from top to bottom due to increase in atomic radii.

Some increase in 1st ionization enthalpy from Si to Ge and Sn to Pb is due to poor shielding effect of d and f electrons.

4. Electronegativity More electronegative than group 13 elements. Electro negativity from Si Pb is same.

5. Physical properties All the members are solids, C, Si are non-metals. Gemetalloid, tin and lead are soft metals.

6. Chemical properties (a) Oxidation state: Group 14 elements have four electrons in their outermost shell (ns2, np2). So elements of this group show an oxidation state of +4 and +2. +2 oxidation state is shown by lower members due to non-availability of ns electrons for bonding due to inert pair effect. Sum of 1st four ionization enthalpies is very high.

So compounds in +4 oxidation state are covalent. Carbon & Silicon show an oxidation state of +4. Carbon also

shows a negative oxidation state. Ge, Sn and lead show an oxidation state of both +2 and +4. Tendency to show

+2 oxidation state increases from Ge to Pb.

Ge forms stable compounds in +4 oxidation state and few compounds in +2 oxidation state. Sn forms compounds in both +2 and +4 oxidation states. (Sn in +2 oxidation state acts as a reducing agent)



In Pb +2 oxidation state is stable and so compounds in +4 oxidation state behave as oxidizing agent. Compounds in +4 oxidation states are electron precise so do not act as electron doner or electron acceptor.

Except carbon other members of the group also show still higher oxidation states by expanding its valence shell

due to presence of nd sub-shell. So their halides undergo hydrolysis and form complexes like [SiF6]2,

[GeCl6]2.

(i) Reaction with Oxygen: All members of the group form oxides and dioxides MO and MO2. SiO exists at high temperature. Oxides in higher oxidation state are more acidic than those in lower oxidation

state. Acidic nature of oxides decreases and basic nature increases from top to bottom. CO2, SiO2,

GeO2 are acidic in nature. SnO2 and PbO2 are amphoteric. CO is neutral. GeO Acidic, SnO and PbO are amphoteric.

(ii) Reactivity with water: Carbon, Silicon and Germanium not effected by water. Tin reacts with steam. Sn + 2H2O SnO2 + 2H2 Lead is unaffected due to formation of protective oxide layer.

(iii) Reaction with halogens: They form halides of the type MX2 and MX4 (where X = F, Cl, Br, I). Except carbon all other members react directly with halogens. Most of the halides in +4 oxidation

states are covalent and the metal atom is sp3 hybridised and the molecules have tetrahedral shape. SnF4

and PbF4 are ionic. PbI4 does not exist. Ge, Sn and Pb also form dihalides. Stability of dihalides

increases from Ge to Pb.

GeX4 is more stable than GeX2

PbX2 is more stable than PbX4

Except CCl4 other tetrahalides are easily hydrolysed by water as the central atom can easily accommodate lone pair

of electrons on O atom of water (due to presence of nd sub-shell). This can be explained by the example of SiCl4.

ANOMALOUS BEHAVIOUR OF CARBON

Carbon differs from other members of the group due to small size, high electronegativity, higher ionization enthalpy

and non-availability of nd sub-shell.

1. Carbon can show a maximum covalency of four other members can expand their covalence due to presence of their nd sub-shell.

2. Due to small size and high electronegativity C has an ability to form p p bonds with itself and with

other atoms of small size (C C, C C, C O, C N, C S). Other members of the group do not form

p p bonds. 3. Carbon has highest tendency of catenation due to

(i) small (ii) strong C-C bond (iii) Tendency of C to show sp, sp2 and sp3 hybridization. (iv) Tendency to form p p bonds.

Carbon shows allotropic forms.

ALLOTROPES OF CARBON

Crystalline Allotropies

1. Diamond



2. Graphite 3. Fullerenes

DIAMOND

In diamond each C atom is sp3 hybridised and linked with four other carbon atoms tetrahedrally forming a 3D- net work solid with C C bond length 154 pm.

Diamond is hardest substance so used as an abrasive and drilling of Rocks. Non-conductor of electricity and good

conductor of heat. It has high refractive index so used as a precious stone.

GRAPHITE

In graphite each carbon atom is sp2 hybridised and linked with three other carbon atoms forming a layer structure

which is composed of hexagonal rings of carbon atoms C C bond length with in the layer is 141.5 pm. The different layers are held by Vander Waals forces with a distance of 340 pm. The fourth electron of each C atom form a cloud of delocalized electrons over the whole sheet. Because of these mobile delocalized electrons graphite conducts electricity. When force is applied one layer of hexagons can easily slide over the other layer. So

graphite is very soft and slippery. So graphite is used as a solid lubricant.

FULLERENES

Fullerenes are made by heating graphite in an electric arc in presence of inert gas like helium and argon. The sooty

material obtained on condensation has C60 with smaller quantities of C70. Fullerenes are cage like molecules. A C60

molecule has a shape like soccer ball with 20, six membered rings and 12, five membered rings. A six membered ring is fused with six or five membered ring, but a five membered ring can only fuse with six membered ring

structures. All the C atoms are sp2 hybridised and linked with three other carbon atoms and remaining electron at



each carbon atom is delocalized which gives aeromatic character. The ball shaped molecule has 60 vertices each one

being occupied by one C atom containing both single and double bonds. C C distances 143.5 pm and 138.3 pm respectively. Such fullerene is known as Buck minster fullerenes or Bucky ball.

Graphite is thermodynamically most stable allotrope of carbon. f Ho of graphite is taken as zero. f Ho for diamond is 1.90 kJ/mole1 and that of C60 is 38.1 kJ mol

1.

Carbon Black By burning hydrocarbons in limited supply of air. Charcoal & Coke By heating wood or coal at high temperature in absence of air.

Uses of Carbon

1. Graphite fibres embedded in plastic material is used in making tennis rackets, fishing rods. 2. Graphite is used as electrodes. 3. Graphite in making crucibles. 4. Activated charcoal is used in gas masks. 5. Activated charcoal is used in water filters to remove organic impurities. 6. Activated charcoal is used in air-conditioning units to control odour. 7. Carbon black in blank ink and filler in automobiles tyres. 8. Coke is used as a fuel and a reducing agent in metallurgy. 9. Diamond is used in jewellery as a precious stone. Measured in carats (1 carat = 200 mg)

IMPORTANT COMPOUNDS OF CARBON & SILICON

CARBON MONOXIDE (CO)

1. Heating carbon in limited supply of air. 2C (s) + O2 (g) 2CO (g)

2. By dehydration of formic acid with conc. H2SO4 at 370 K. H.COOH H2O + CO

3. By passing steam over red hot coke.

C (s) + H2O (g) K1273473

CO (g) + H2 [Water or synthesis gas]

4. By passing air over red hot coke.

2C (s) + O2 (g) + 4N2 K1273

2CO (g) + 4N2 (g) [Producer gas]

Properties



Colourless, odourless gas insoluble in water, neutral in nature. It combines with haemoglobin to give a stable

carboxy haemoglobin complex. This decreases oxygen carrying capacity of haemoglobin ultimately resulting into

death. Due to presence of a lone pair of electrons it reacts with metals to give metal carboxyl.

Uses

1. Water gas & producer gas are used as industrial fuels. 2. As a strong reducing agent. Reduces metal oxides to metals.

Fe2O3 (s) + 3CO (g)

2Fe (s) + 3CO2 (g)

ZnO (s) + CO (g)

Zn (s) + CO2 (g)

CARBON DIOXIDE (CO2)

1. By complete combustion of carbon and carbon containing fuels.

C (s) + O2 (g)

CO2 (g)

CH4 (g) + 2O2 (g)

CO2 (g) + 2H2O (l)

2. By the actions of dilute acids on lime stone. CaCO3 (s) + 2HCl (aq) CaCl2 (aq) + CO2 (g) + H2O (l)

3. By heating lime stone CaCO3 (s) CaO (s) + CO2 (g)

Properties

1. It is colourless, odourless gas. Sparingly soluble in water. Dissolves in water to give carbonic acid. H2O (l) + CO2 (g) H2CO3 (aq) Carbonic acid is a dibasic acid. H2CO3 (aq) + H2O (l) HCO3

(aq) + H3O+ (aq)

HCO3 (aq) + H2O (l) CO32 (aq) + H3O

+ (aq)

HCO3 / H2CO3 buffer maintenance the pH of blood between 7.26 to 7.42.

2. Acidic in nature so combines with alkalies to give metal carbonates. 2NaOH (aq) + CO2 (g) Na2CO3 (aq) + H2O (l)

3. Green pigments of the plants (Chlorophyll) converts CO2 of atmosphere to carbohydrates, this phenomenon is known as Photosynthesis.

6CO2 + 6H2O lChlorophyl

C6H12O6 + 6O2

4. It is non-toxic in nature. 5. Increase in contents of CO2 in the atmosphere lead to green house effect. Leading to give rise in global

temperature.

6. Solid CO2 is known as dry ice which is used as refrigerant for ice cream and frozen food. 7. Gaseous CO2 is used in soft drinks. 8. As a fire extinguisher. 9. In manufacture of urea.

Structure of CO2 C atom is sp hybridized. The hybrid orbitals overlap with P atomic orbital of oxygen atoms to give bonds. The two electrons of C atom are involved in p p bonding with oxygen. This results in linear shape with both C C bond of equal length (115 pm).

RESONANCE STRUCTURE

Carbon dioxide is a resonance hybrid of following structure

SILICON DIOXIDE (SIO2)



Silicon dioxide (Silica) occurs in several crystallographic forms Crystalline forms of silica are Quartz, Cristobalite, Tridymite. These are inter convertible at suitable temperature. Silicon dioxide is a covalent, 3D net

work solid. Each silicon atom is tetrahedrally linked with four oxygen and each O atom with two silicon atoms.

Silica is almost non-reactive due to very high SiO bond enthalpy. Not attacked by halogens, H2, acids and metals. Attacked by HF and NaOH.

SiO2 + 2NaOH Na2SiO3 + H2O Sodium silicate

SiO2 + 4HF SiF4 + 2H2O SiF4 + 2HF H2SiF6 (Fluoro silicic acid) soluble in water.

Uses of Silica

1. Quartz is used as piezoelectric material in accurate clocks, modern radio and television broadcasting and mobile radio communications.

2. Silica gel is used as a drying agent and in chromatography. 3. Amorphous silica is used in filteration plants.

SILICONES

These are organ silicon polymers with (R2 Si O ) as a repeating unit. They are prepared from alkyl or aryl substituted silicon chlorides. Hydrolysis of dimethyl-dichloro silane (CH3)2CCl2 followed by condensation

polymerization give a straight chain polymer. The chain length of polymer can be controlled by adding (CH3)3CCl.



Silicones are water repelling in nature. High thermal stability, high dielectric strength and resistance to oxidation

and chemicals. They are used as sealants, greases, electrical insulators, water proofing of fabrics. In surgical and

cosmetic plants.

SILICATES

Common silicates found in nature are feldspar, zeolites, mica and Asbestos. Basic structural units of silicate are SiO4

4.

Silicate units may be joined together via corners by sharing 1, 2, 3 or 4 oxygen atoms per silicate units to give chain,

ring, sheet or 3-D structures. Important man made silicates are glass and cement.

Zeolites

If Alluminium atoms replace few silicon atoms in 3D network of SiO2 overall structure is known as aluminosilicate.

It acquires a negative charge. Cations such as Na+, K+, Ca2+ balance the negative charge. Examples feldspar and

zeolitees. Zeolites are widely used as a catalyst in photochemical industries for checking of hydrocarbons and

isomerisation.

Z S M -5 used to convert alcohols directly into gasoline.

Hydrated zeolites are used as ion exchangers.

Tin shows two allotropic forms white tin which is crystalline in nature and grey tin which is amorphous (i) Tin plague : Due to conversion of crystalline white tin two amorphous grey tin, so crumbling of tin

takes place. This is known as tin plague

(ii) Tin cry : on bending a tin seat there is a peculiar sound which is known as tin cry. This is due to rubbing of the crystals of tin.

p block element

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