Chapter 11 s & p Block Elements part 2 by TEACHING CARE Online coaching and tuition classes

Chapter 11 s & p Block Elements part 2 by TEACHING CARE Online coaching and tuition classes

File name : Chapter-11-s-p-Block-Elements-part-2.pdf

 

 

2
  • Gypsum (CaSO4 . 2H 2O) : On partially dehydrates to produce plaster of

 

 

CaSO

. 2H O

120 o C

1       +   1

 

4

Gypsum

2   ¾¾¾¾® CaSO4 .    H 2O

Plaster of paris

1 2 H 2 O

 

 

 

Plaster of paris : CaSO4

1

. 2 H

O ¾¾HO ® CaSO

2
4

Setting

. 2H

2 O ¾¾Ha¾rde¾nin¾g  ®

CaSO4

.2H 2 O

 

Plaster of paris

 

4

Gypsum ¾¾200¾o¾C ® CaSO

orthorhombic

 

(anhydrous)

Monoclinic (gypsum)

 

(dead burnt plaster)

 

 

Gypsum when heated to about

200o C

is converted into anhydrous calcium sulphate. The anhydrous form

 

(anhydrite) is known as dead burnt plaster because it does not set like plaster of paris when moistened with water.

 

  • Calcium Hydroxide

Ca(OH)2 (slaked lime)

 

 

CaO + H 2 O ® Ca(OH)2 ;  Ca(OH)2  + CO2  ®CaCO3  + Ca(HCO3 )2

Suspension of Ca(OH)2 in water is called milk of lime.

Ca(OH)2 + Cl2 ® CaOCl2 + H 2 O

Group 13 of long form of periodic table (previously reported as group III A according to Mendeleefs periodic table) includes boron (B) ; aluminium (Al) , gallium (Ga), indium (In) and thallium (Tl) Boron is the first member of group 13 of the periodic table and is the only non-metal of this group. The all other members are metals. The non- metallic nature of boron is due its small size and high ionisation energy. The members of this family are collectively known as boron family and sometimes as aluminium family.

Electronic configuration

 

Element      Electronic configuration (ns 2 np1)
5 B                     1s 2 ,2s 2 2p1 or [He] 2s 2 2p1

13 Al                   1s 2 ,2s 2 2p6 ,3s 2 3p1 or [Ne] 3s 2 3p1

31 Ga                  1s 2 ,2s 2 2p6 ,3s 2 3p6 3d10 ,4s 2 4 p1 or [Ar]3d10 4s 2 4 p1

49 In                  1s 2 ,2s 2 2p6 ,3s 2 3p6 3d10 ,4s 2 4 p6 4d10 ,5s 2 5p1 or [Kr]4d10 5s 2 5p1

81 Tl                   1s 2 ,2s 2 2p6 ,3s 2 3p6 3d10 ,4s 2 4 p6 4d10 4 f 14 ,5s 2 5p6 5d10 ,6s 2 6p1 or [Xe]4 f 14 5d10 6s 2 6p1

Physical properties

  • A regular increasing trend in density down the group is due to increase in
  • Melting points do not vary regularly and decrease from B to Ga and then
  • Boron has very high pt because it exist as giant covalent polymer in both solid and liquid state.
  • Low pt of Ga (29.80C) is due to the fact that consists of only Ga2 molecule; it exist as liquid upto 20000C and hence used in high temperature thermometry.

 

 

 

  • Boiling point of these elements however show a regular decrease down the
  • The abrupt increase in the atomic radius of Al is due to greater screening effect in Al (it has 8 electrons in its penultimate shell ) than in B (it has 2 electrons in its penultimate shell)
  • The atomic radii of group 13 elements are smaller than the corresponding s-block elements. This is due to the fact that when we move along the period, the new incoming electron occupy the same shell whereas the nuclear charge increases regularly showing more effective pull of nucleus towards shell This ultimately reduces the atomic size.
  • The atomic radius of Ga is slightly lesser than of Al because in going from Al to Ga, the electrons have already occupied 3d sub shell in The screening effect of these intervening electrons being poor and has less influence to decrease the effective nuclear charge, therefore the electrons in Ga experience more forces of attractions towards nucleus to result in lower size of Ga than Al

(9)    Oxidation state

  • All exhibit +3 oxidation state and thus complete their octet either by covalent or ionic
  • Boron being smaller in size cannot lose its valence electrons to form B3+ ion and it usually show +3 The tendency to show +3 covalence however decreases down the group Even Al shows +3 covalence in most of its compounds.
  • Lower elements also show +1 ionic state g Tl +, Ga+. This is due to inert pair effect. The phenomenon in which outer shell ‘s’ electrons (ns2) penetrate to (n-1) d electrons and thus become closer to nucleus and are more effectively pulled the nucleus. This results in less availability of ns2 electrons pair for bonding or ns2 electron pair becomes inert. The inert pair effect begins after n ³ 4 and increases with increasing value of n.
  • The tendency to form M+ ion increases down the Ga+1 < Tl+1
  • Hydrated ions : All metal ions exist in hydrated

(11)    Ionisation energy

  • Inspite of the more charge in nucleus and small size, the first ionisation energies of this group elements are lesser than the corresponding elements of s block. This is due to the fact that removal of electron from a p-orbitals (being far away from nucleus and thus less effectively held than s-orbitals) is relatively easier than s-orbitals.
  • The ionisation energy of this group element decrease down the group due to increases in size like other group
  • However, ionisation energy of Ga are higher (table ) than that of Al because of smaller atomic size of Ga due to less effective shielding of 3d electrons in Ga. Thus valence shell exert more effective nuclear charge in Ga to show higher ionisation

(12)    Electropositive character

  • Electropositive character increases from B to Tl.
  • Boron is semi metal, more closer to non-metallic nature whereas rest all members are pure
  • Furthermore, these elements are less electropositive than s-block elements because of smaller size and higher ionisation

 

 

 

 

 

(13)    Oxidation potential

  • The standard oxidation potentials of these element are quite high and are given below,

 

 BAlGaInTl
E0op for M ® M3++ 3e+1.66+0.56+0.34+1.26
E0op for M ® M+ + e+0.55+0.18+0.34
  • However Boron does not form positive ions in aqueous solution and has very low oxidation
  • The higher values of standard oxidation potentials are due to higher heats of hydration on account of smaller size of trivalent
  • Aluminium is a strong reducing agent and can reduce oxides which are not reduced even by This is due to lower ionisation energy of aluminium than carbon. The reducing character of these elements is Al > Ga > In > Tl.
    • Complex formation : On account of their smaller size and more effective nuclear charge as well as vacant orbitals to accept elements, these elements have more tendency to form complexes than-s block

Chemical properties

  • Occurrence : The important of this group elements are given below,

Boron : Borax (Tincal)             (Na2B4O7.10H2O), Colemanite         (Ca2B6O115H2O) Boracite                             (2Mg3B8O15.MgCl2), Boronatro calcite (CaB4O7.NaBO2.8H2 O), Kernite               (Na2B4O7.4H2O), Boric acid             (H3BO3)

Aluminium : Corundum (Al2O3), Diaspore (Al2O3.H2O), Bauxite (Al2O3. 2H2O), and Cryolite (Na3AlF6).

(2)  Hydrides

  • Elements of gp 13 do not react directly with hydrogen but a number of polymeric hydrides are known to

exist.

  • Boron forms a large of volatile covalent hydrides, known as boranes e.g. B2 H6,B4H10,B5H11 ,B6H10

Two series of borones with general formula BnHn + 4 and BnHn + 6 are more important.

  • Boranes are electron deficient It is important to note that although BX3 are well known, BH3 is not known. This is due of the fact that hydrogen atoms in BH3 have no free electrons to form pp-pp back bonding and thus boron has incomplete octet and hence BH3 molecules dimerise to form B6H6 having covalent and three centre bonds.
  • Al forms only one polymeric hydride (AlH3)n commonly known as alane It contains A1…..H……Al
  • Al and Ga forms anionic hydrides e. LiAlH4 and Li Ga H4 , 4 LiH + AlCl3 ¾¾eth¾er ® Li[AlH4 ] + 3LiCl

(3)  Reactivity towards air

  • Pure boron is almost unreactive at ordinary It reacts with air to form B2O3 when heated It does react with water. Al burns in air with evolution of heat give Al2O3.
  • Ga and In are not effected by air even when heated whereas Tl is little more reactive and also form an oxide film at In moist air , a layer of Tl (OH) is formed.

 

 

 

  • Al decomposes H2O and reacts readily in air at ordinary temperature to form a protective film of its oxides which protects it from further

(4)  Oxides and hydroxides

  • The members of boron family form oxide and hydroxides of the general formula M2O3 and M (OH)3
  • The acidic nature of oxides and hydroxides changes from acidic to basic through amphoteric from B to

B2O3 and B(OH)3> Al2O3 and Al(OH)3 >Ga2O3 and Ga(OH)3> In2O3 In (OH)3> Tl2O3 Tl(OH)3

(acidic)          (amphoteric)          (amphoteric)           (basic)          (strong basic)

Note : ® B(OH)3 or H3BO3 is weak monobasic Lewis acid.

  • Boric acid, B(OH)3 is soluble in water as it accepts as it accepts lone pair of electron to act as Lewis Rest all hydroxides of group 13 are insoluble in water and form a gelatinous precipitate.

B(OH)3 + H2O ®B(OH)41-+H+

  • Al2O3 being amphoteric dissolves in acid and alkalies

 

Al2O3 + 3H2SO4® Al2 (SO4)3 + 3H2O   ;

Al 2 O3  + 2NaOH ¾¾fu¾se ®

2NaAlO3

Sodium meta aluminate

  • H 2 O

 

  • One of the crystalline form of alumina (Al2O3) is called It is very hard and used as abrasive. It is prepared by heating amorphous form of Al2O3 to 2000 K.

(5)  Action of Acids

  • Boron does not react with non oxidizing acids, however, it dissolves in nitric acid to form boric
  • Al, Ga and In dissolve in acids forming their trivalent cations; however, Al and Ga become passive due to the formation of protective film of
  • Thallium dissolves in acids forming univalent cation and becomes passive in HCl due to the formation of water insoluble

(6)  Action of Alkalies

  • Boron dissolves only in fused alkalis, 2B + 6NaOH (fused)®2Na3BO3 + 3H2
  • Al and Ga dissolves in fused as well as in aqueous alkalis, 2Al + 2 NaOH + 2H2O ®2NAl O2 + 3H2
  • Indium remains unaffected in alkalies even on

(7)  Halides

  • All the group 13 elements from the trihalides, MX3 on directly combining with

M + X2 ® MX3

  • All the trihalides of group 13 elements are known except Tl (III)
  • Due to small size and high electronegativity of boron, all boron halides are covalent and Lewis These exist as monomeric molecules having plane triangular geometry (sp2 hybridization).
  • All Boron trihalides except BF3 are hydrolysed to boric BX3+ 3H2O ®B(OH)3 + 3HX; [X=Cl,Br,I]

 

However, BF3 forms as addition product with water,BF3 + H2O®H+ [BF3OH]

H2O

H3O+ [BF3OH] .

 

 

 

BF3 having less tendency for hydrolysis as well as Lewis acid nature, is extensively used as a catalyst in organic reactions e.g. Friedel- Crafts reaction.

  • Boron atom, in BX3, has six electrons in the outermost orbit and thus it can accept a pair of electrons form a donor molecule like NH3 to complete its Hence boron halides act as very efficient Lewis acids. The relative Lewis acid character of boron trihalides is found to obey the order ; BI3 > BBr 3 > BCl3 > BF3.

However, the above order is just the reverse of normally expected order on the basis relative electronegativities of the halogens. Fluorine, being the most electronegative, should create the greatest electron deficiency on boron and thus B in BF3 should accept electron pair from a donor very rapidly than in other boron trihalides. But this is not true.

This anomalous behaviour has been explained on the basis of the relative tendency of the halogen atom to back-donate its unutilised electrons to the vacant p orbitals of boron atom. In boron trifluoride, each fluorine has completely filled unutilised 2p orbitals while boron has a vacant 2p orbital. Now since both of these orbitals belong to same energy level (2p) they can overlap effectively as a result of which fluorine electrons are transferred into the vacant 2p orbital of boron resulting in the formation of an additional pp – pp bond. This type of bond formation is known as back bonding or back donation. Thus the B- F bond has some double bond character. Back bonding may take place between boron and of the three fluorine atoms and thus boron trifluoride is regarded as a resonance hybrid of some structures.

Resonance in boron trifluoride is also evidenced by the fact that the three boron-fluorine bonds are indentical and are shorter than the usual single boron-fluorine bond As a result of back bonding, the electron deficiency of boron is reduced and hence Lewis acid nature is decreased. The tendency for the formation of back bonding (pp- pp bond) is maximum in BF3 and decreases very rapidly from BF3 to BI3 This is probably due to the fact that overlapping of the vacant 2p orbitals of boron cannot take place easily with the p-orbitals of high energy levels (3p in Cl, 4p in Br and 5p in iodine). Thus BI3 Br3 and BCl3 are stronger Lewis acids than the BF3.

  • Lewis acid character of halides of the group 13 elements decreases in the order,B > Al > Ga > In
  • Boron halides form complex halides of the type, [BF4], in which boron atom extends its coordination number to four by utilising empty p-orbital. It cannot extend its coordination number beyond four due to non availability of d-orbitals. However, the other trihalides of this group form complex halides of the type (AlF6)3-, (GaCl6)3- and (InCl6)3-, etc where the central atom extends its coordination number to 6 by the use of d-orbitals.
  • The fluorides of Al, Ga In and Tl are ionic and have high melting points. The high melting points of metal fluorides can be explained on the basis that their cations are sufficiently large and have vacant d-orbitals for attaining a coordination number of six towards the relatively small fluorine
  • Other halides of Al, Ga, In and Tl are largely covalent in anhydrous state and possess low pt. These halides do not show backbonding because of increases in the size of the element. However, the make use of vacant p-orbitals by co-ordinate bond i.e. metal atoms complete their octet by forming dimers. Thus aluminium chloride, aluminium bromide and indium iodide exist as dimers, both in the vapour state and in non-polar solvents.

The dimer structure for Al2Cl6 is evidenced by the following facts,

  • Vapour density of aluminium chloride measured at 4000C corresponds to the formula Al2 Cl6.
  • Bond distance between aluminium chlorine bond forming bridge is greater (2.21A0) than the distance between aluminum-chlorine bond present in the end (2.06 A0). The dimeric structure disappears when the halides

 

 

 

are dissolved in water This is due to high heat of hydration which split the dimeric structure into [M(H2O)6]3+ and 3X ions and the solution becomes good conductor of electricity.

Al2Cl6 + 2H2O ®2[Al(H2O)6]3++6Cl ; Therefore Al2Cl6 is ionic in water.

The dimeric structure may also split by reaction with donor molecules e.g. R3N. This is due to the formation of complexes of the type R3NAl Cl3 The dimeric structure of Al2 Cl6 exist in vapour state below 473 K and at higher temperature it dissociates to trigonal planar AlCl3 molecule.

Note : ® Boron halides do not exist as dimer due to small size of boron atom which makes it unable to co-ordinate four large-sized halide ions.

  • BF3 and AlCl3 acts as catalyst and Lewis acid in many of the industrial

Anomalous Behaviour of Boron

Like Li and Be, Boron – the first member of group 13 also shows anomalous behaviour due to extremely low size and high nuclear charge/size ratio, high electronegativity and non-availability of d electrons. The main point of differences are,

  • Boron is a typical non- metal whereas other members are
  • Boron is a bad conductor of electricity whereas other metals are good
  • Boron shows allotropy and exists in two forms – crystalline and Aluminium is a soft metal and does not exist in different forms.
  • Like other non-metals, the melting point and boiling point of boron are much higher than those of other elements of group
  • Boron forms only covalent compounds whereas aluminium and other elements of group 13 form even some ionic
  • The hydroxides and oxides of boron are acidic in nature whereas those of others are amphoteric and

basic.

  • The trihalides of boron (BX3) exist as monomers On the other hand, aluminium halides exist as dimers (Al2X6).
  • The hydrides of boron e. boranes are quite stable while those of aluminium are unstable.
  • Dilute acids have no action on boron Others liberate H2 from
  • Borates are more stable than
  • Boron exhibit maximum covalency of four g., BH4 ion while other members exhibit a maximum

covalency of six e.g., [Al(OH)6]3-.

  • Boron does not decompose steam while other members do
  • Boron combines with metals to give borides g. Mg3 B2. Other members form simply alloys.
  • Concentrated nitric acid oxidises boron to boric acid but no such action is noticed other group B + 3HNO3 ®H3BO3 + 3NO2

Diagonal relationship between Boron and Silicon

Due to its small size and similar charge/mass ratio, boron differs from other group 13 members, but it resembles closely with silicon, the second element of group 14 to exhibit diagonal relationship. Some important similarities between boron and silicon are given below,

 

 

 

  • Both boron and silicon are typical non-metals, having high pt. b.pt nearly same densities (B=2.35g ml –1 S=2.34 g//ml). low atomic volumes and bad conductor of current. However both are used as semiconductors.
  • Both of them do not form cation and form only covalent
  • Both exists in amorphous and crystalline state and exhibit
  • Both possess closer electronegativity values (B=2.0;Si=1.8).
  • Both form numerous volatile hydrides which spontaneously catch fire on exposure to air and are easily
  • The chlorides of both are liquid, fume in most air and readily hydrolysed by

BCl3 + 3H2O ® B(OH)3 + 3HCl ; SiCl4 + H2O ®Si(OH)4 + 4HCl

  • Both form weak acids like H3BO3 and H2SiO3.
  • Both form binary compounds with several metals to give borides and These borides and silicide react with H3PO4 to give mixture of boranes and silanes.

3Mg + 2B ® Mg3B2 ; Mg3B2 + H3PO4 ® Mixture of boranes

(Magnesium boride)

 

2Mg + Si ® Mg2Si ; Mg2Si + H3PO4 ®Mixture of silanes

(magnesium silicide)

 

  • The carbides of both Boron and silicon (B4 C and SiC) are very hard and used as
  • Oxides of both are acidic and can be reduced by limited amount of Mg In excess of Mg boride and silicide are

B2O3 + 3Mg ® 3MgO + 2B ; SiO2 + 2Mg ® 2MgO + Si

  • Both the metals and their oxides are readily soluble in

2B + 6NaOH ® 2Na3BO3 + 3H2 ­ ;  Si + 2NaOH + H2O ® Na2SiO3 + 2H2 ­

(borate)                                                                                     (silicate)

 

B2O3 + 6NaOH ® 2Na3BO3 + 3H2O ; SiO2 + 2NaOH ® Na2SiO3 + H2O

 

Both borates and silicates have tetrahedral structural units

BOn and

SiO n respectively. Boro silicates are

 

4
4

known in which boron replaces silicon in the three dimensional lattice. Boron can however form planar BO3 units.

  • Acids of both these elements form volatile esters on heating with alcohol in presence of conc. H2SO4. B(OH)3 + 3ROH®B(OR)3 + 3H2O ; Si(OH)4 + 4ROH® Si(OR)4 + 4H2O

Boron and its compounds

Boron is the first member of group –13 (IIIA) of the periodic table. Boron is a non- metal . It has a small size

 

and high ionization energy due to which it can not lose its valence electrons to form especially the hydrides and halides are electron deficient and behave as Lewis acid.

(1)  Ores of boron

  • Borax or tincal : Na2 B4O7 . 10 H2O
  • Kernit or Rasorite : Na2 B4O7 . 10 H2O

B +3

ion. Its compounds

 

 

 

  • Colemanite : Ca2 B6O11 . 5 H2O
  • Orthoboric acid : H3BO3 (It occurs in the jets of steam called soffioni escaping from ground in the volcanic region of the Tuscany). Boron is present to a very small extent (0.001%) in earth’s
  • Isolation : Elemental boron in the form of dark brown powder is obtained either by reduction of boric oxide with highly electropositive metals like K, Mg, Al, Na, in the absence of air and boron halides with

 

hydrogen at high temperature eg. B2O3 + 6K

¾¾He¾at ® 2B + 3K2O;      2BCl3 + 3H2

¾¾127¾0¾K ® 2B + 6HCl.

 

By thermal decomposition of boron tri-iodide over red hot tungsten filament and boron hydrides for example,

 

  • BI3

¾¾W,h¾e¾at ® 2B + 3I2 ; B2H6

¾¾He¾at ® 2B + 3H2

 

  • Properties : It exists in mainly two allotropic forms e. amorphous dark brown powder and crystalline

 

black very hard solid. It occurs in two isotopic forms, i.e.,

B10

(20% abundance) and

5 B11

(80% abundance).

 

With air, boron forms

B2O3

and BN at 973K, with halogens, trihalides (BX3 )

are fromed with metals, borides are

 

formed. eg.      4B+ 3O2

¾¾He¾at  ®

2B2 O3

Boron trioxide

; 2B +

N2  ¾¾He¾at  ®

2BN     ;

Boron nitride

 

2B + 3X2 ¾¾®

2BX3

Boron trihalide

;        3Mg + 2B

¾¾He¾at  ®

Mg3 B2

Magnesium boride

 

Water, steam and HCl have no action on B. oxidising acids (HNO3 ,

H 2 SO4 ) convert boron to

H3 BO3 .

 

B + 3 HNO3

¾¾®

H3 BO3 + 3NO2 ;  2B +  3 H 2 SO4

¾® 2 H3BO3 + 3 SO2

 

Fused alkalies ( NaOH , KOH ) dissolve boron forming borates, liberating hydrogen.

 

2B + 6KOH

  • F¾u¾sed ® 2 K3 BO3

+ 3H 2

 

  • Uses of Boron : Boron is used in atomic reactors as protective shields and control rods, as a semiconductors for making electronic devices in steel industry for increasing the hardness of steel and in making light composite materials for air

(5)  Compounds of Boron

Boron Hydrides

 

Boron forms hydrides of the types

Bn Hn+4

and

Bn Hn+6

called boranes. Diborane is the simplest boron

 

hydride which is a dimer of

BH 3 .

 

Structure of diborane :

B2 H6

has a three centre electon pair bond also called a banana shape bond.

 

(i)

B Ht

: It is a normal covalent bond (two centre electron pair bond i.e., 2c – 2e).

 

(ii)

B Hb : This is a bond between three atoms,

B Hb B, (three centre electron pair bond i.e., 3c – 2e).

 

 

 

 

 

 

The other boron hydrides are

B5 H9 , B4 H10 , B5 H11 etc.

Boron Halides

 

Boron reacts with halogens on strong heating to form boron halides .

2B + 3 X 2  ¾¾He¾at ® 2B X3 (X  = F, Cl, Br, I)

 

BF3 and

BCl3 are gases,

BBr3

is a volatile liquid while

BI3

is a solid.

 

In these halides, the central boron atom has three shared pairs of electrons with the halogen atoms. Therefore, these have two electrons less than the octet and are electron deficient compounds. They acts as Lewis acids.

F            H                                 F               H

|              |                                   |                 |

FB   +   : NH   ¾¾® FB ¬¾¾ NH

 

|

F

Lewis acid

|                                   |                 |

H                                F               H

Lewis base

 

The relative acidic strength of boron trihalides decreases as :

BI 3 > BBr3 > BCl 3 > BF3 .

 

Borax ( Na2 B4 O7 .10H 2 O )

It occurs naturally as tincal (Suhaga) which contains about 50% borax in certain land, lakes. It is also

 

obtained from the mineral colemanite by boiling it with a solution of

Na2CO3 .

 

Ca2 B6 O11 + 2Na2 CO3 ¾¾® Na2 B4 O7 + 2CaCO3  + 2NaBO2

Colemanite                                                 Borax

Properties : (i) Its aqueous solution is alkaline due to hydrolysis,

 

Na2  B4 O7 + 7HO

¾¾®

2NaOH+4 H 3 BO3 .

 

  • On heating borax loses its water of crystallization and swells up to form a fluffy mass. On further heating, it melts to give a clear liquid which solidifies to a transparent glassy bead consisting of sodium metaborate ( NaBO2 )

 

and boric anhydride ( B2O3 ),     Na2 B4 O7 .10H2O ¾¾D ® Na2 B4 O7 ¾¾D ® 2NaBO2 +B2 O3 .

 

-10 H 2 O

Borax bead

 

Borax bead is used for the detection of coloured basic radicals under the name borax bead test.

 

  • When heated with with a green edged

C2 H 5 OH and conc.

H2SO4

it gives volatile vapours of triethyl borate which burns

 

Na2B4O7 + H2SO4 + 5H2O ¾¾® Na2SO4 + 4 H3 BO3 ;

H3 BO3 + 3C2 H5 OH ¾¾® B(OC2 H5 )3 + 3H2O

Triethyl borate

 

This reaction is used as a test for borate radical in qualitative analysis.

Uses : (1) In making optical and hard glasses. (2) In the laboratory for borax bead test. (3) In softening of water. (4) In the preparation of medicinal soaps due to its antiseptic character.

 

Borax bead test : Borax bead is a mixture of

NaBO2 and

B2O3 .

B2O3

on heating combines readily with a

 

number of coloured transition metal oxides such as Co, Ni, Cr, Cu, Mn, etc. to form the corresponding metaborates

 

which possess characteristic colours,

CoSO4 ¾¾D ® CoO + SO3 ;

CoO + B2O3 ¾¾®

Co(BO2 )2

Cobalt meta borate (Blue)

 

 

 

Colours of some important metaborates are: Cupric metaborate, Cu(BO 2 )2 is dark blue, chromium metaborate,

 

Cr (BO2 )2 is green, nickel metaborate,

Ni(BO2 )2 is brown and manganese metaborate,

Mn(BO2 )2

is pink violet.

 

Boric acid or orthoboric acid ( H 3 BO3 )

 

It is obtained from borax by treating with dil. HCl or dil.

H2SO4  ,

 

Na2 B4 O7 + 2HCl + 5H 2 O ¾¾® 2NaCl + 4 HBO3

It can also be obtained from the mineral colemanite by passing

SO2

through a mixture of powdered mineral

 

in boiling water,

Ca2 B6 O11 +

4SO2 + 11H 2 O ¾¾® 2Ca(HSO3 )2 + 6HBO3

 

Properties : (i) It is a very weak monobasic acid, does not act as a proton doner but behaves as a Lewis

 

acid i.e. it accepts a pair of electrons from OH ion of

HO ,

H 3 BO3  + H 2 O ¾¾®[B(OH)4 ]  + H +

 

It acts as a strong acid in presence of polyhydroxy compounds such as glycerol, mannitol etc. and can be titrated against strong alkali .

  • With NaOH it forms, sodium metaborate, H 3 BO3 + NaOH ¾¾® NaBO2 + 2H 2 O
  • With C2 H 5 OH and H2SO4 , it gives triethyl borate

H3BO3 + 3 C2 H 5 OH    ¾¾Con¾c. H¾2SO¾4  ® B(OC2 H 5 )3 + 3 H 2 O

  • Action of heat : The complete action of heat on boric acid may be written as,

 

H3 BO3  ¾¾373¾K ®

Boric acid

HBO2

Metaboric acid

¾¾433¾K  ®  H 2 B4 O7

Tetra boric acid

¾¾Red¾h¾ot ®

B2 O3

Boron oxide

 

3

Structure : In boric acid, planar BO -3 units are joined by hydrogen bonds to give a layer structure.

Uses : (i) As a food preservative. (ii) As a mild antiseptic for eye wash under the name boric lotion. (iii) For the preparation of glazes and enamels in pottery.

 

Borazine or Borasole or Triborine triamine (

B3 N 3 H 6 )

 

It is a compound of B, N and H. It is a colourless liquid and is also called inorganic benzene.

2    6                3                             3    3    6               2

2 B  H   + 6 NH   ¾¾180o¾C ® 2B  N  H   + 12H   .

It has a six membered ring of alternating B and N atoms, each is further linked to a H– atom.

Boron nitride (BN)

 

It is prepared by treating

BCl 3

with an excess of

NH 3 and pyrolysing the

 

resulting mixture in an atmosphere of

NH 3 at 750 o C ,

 

 

 

BCl   + NH       ¾® [ H

N ¾¾® BCl

] ¾¾750¾o¾C ® BN + 3HCl .

 

3              3                     3

  • Excess NH3

 

It is a colourless, good insulator, diamagnetic and almost unreactive solid

Aluminium and its compounds

 

  • Ores of Aluminium : Bauxite

(Al 2 O3 .2H 2 O), Cryolite

(Na3 AlF6 , Felspar

(KAlSi3O8 ),

Kaolinite

 

(Al 2 O3 . 2SiO2 .2H 2 O) , Mica

(K2 O.3 Al2 O3 . 6SiO2 .2H 2 O),

Corundum

(Al 2O3 ) , Diaspore

(Al 2 O3 .H 2 O),

Alunite

 

or alum stone [K 2 SO4 . Al2 (SO4 )3 . 4 Al(OH)3 ].

  • Extraction : Aluminium is obtained by the electrolysis of the oxide (alumina) dissolved in fused This involves following steps,

 

 

 

Purification of ore

  • Baeyer’s process

 

Bauxite

  • R¾oas¾ted ®Roastedore ¾+¾Ca¾ustic¾soda¾solu¾t¾ion ®

Filtrate

  • F¾ilte¾red ® Pure Al O
  • H¾eat® Al(OH )

 

Finely powdered (red)

FeO® Fe2O3

High pressure (150o C, 80 atm) filtered, Fe2O3 as residue

(Sod. Aluminate)

CO2                                       2     3                                                        3

 

  • Hall’s process

 

Bauxite    ¾¾+N¾a2C¾O3¾® Solution ¾¾C¾O2¾¾® Precipitate Al(OH)

¾¾He¾at ® Pure  Al  O

 

(Finely powdered) Fused, extracted with

50-60o C and filtered.

3                                     2    3

 

(red)

water. Residue Fe2O3

Filtrate (Na 2CO3 )

 

  • Serpek’s process

 

Bauxite

(Finely powdered)

(white)

¾¾+C¾oke¾+N¾2¾®

Heated to

1800o C

Silica reduced to Si which volatalises

  • lumina form aluminium nitride

AlN ¾¾Hyd¾roly¾¾sis ® Pure   Al 2 O3 ¾¾Hea¾t¾ed ® Al(OH)3

 

  • Hall and Heroult process : It is used for extraction of In this process a fused mixture of alumina (20%), cryolite (60%) and fluorspar (20%) is electrolysed using carbon electrodes. Whereas cryolite makes Al 2O3 conducting fluorspar decreases the m.pt. of alumina.

Note : ® Aluminium is refined by Hoope’s electrolytic process.

(3)  Compounds of Aluminium

 

  • Aluminium oxide or Alumina

(Al 2 O3 )

: It occurs in nature as colourless corundum and several

 

coloured minerals like ruby (red), topaz (yellow), Sapphire (blue), amethyst (violet) and emerald (green). These minerals are used as precious stones (gems).

  • Aluminium chloride (Al 2 Cl6 ) : It is prepared by passing dry chlorine over aluminium

Al2 O3  + 3C + 3Cl2  ®  2AlCl3   + 3CO(g)

 

(anhydrous)

It exists as dimer Al 2Cl6 , in inert organic solvents and in vapour state. It sublimes at 100 o C

under vacuum.

 

Dimeric structure disappears when when exposed to air.

AlCl 3

is dissolved in water. It is hygroscopic in nature and absorbs moisture

 

  • Thermite : A mixture of aluminium powder and

FeO3

in the ratio 1:3. It is used for welding of iron.

 

The reaction between Al and

FeO3

is highly exothermic,

Al + Fe2 O3 ® Al 2 O3 + Fe + Heat

 

  • Aluminium sulphate [Al2(SO4)3] : It is used for the preparation of alums g., potash alum

Al 2 (SO4 )3 . KSO4 . 24 HO . It is also used for making fire proof clothes.

Carbon is the first member of group 14 or IVA of the periodic table. It consists of five elements carbon (C), silicon (Si), germanium (Ge), tin (Sn) and lead (Pb). Carbon and silicon are nonmetals, germanium is metalloid and tin and lead are metals.

Electronic configuration

 

Elements        Electronic configuration (ns2 np2)
6 C                         1s 2 ,2s 2 2p 2 or [He] 2s 2 2p 2

14 Si                       1s 2 ,2s 2 2p6 ,3s 2 3p 2 or [Ne] 3s 2 3p 2

 

 

 

Physical properties

  • Non-metallic nature : The non-metallic nature decreases along the

 

CSiGeSnPb
Non-metal metalloidmetalmetal

or semi metal

  • Abundance : Carbon and silicon are most abundant elements in earth’s crust whereas germanium occurs only as traces. Tin and lead also occur in small amounts. Only carbon occurs in free state as coal, diamond and graphite and in combined state as carbonates, CO2 petroleum and natural gas Silicon is the second most abundant element after oxygen in earth’s crust in form of silicates and silica. Germanium found in traces in coal and in certain It important constituent for making conductors and transistors The important ore of tin is tin stone (SnO2) or cassiterite. Lead is found is form of galena (PbS) anglesite (PbSO4) and cerussite (PbCO3) The abundance ratio in earth’s crust is given below,
ElementCSiGsSnPb
Abundance in earth’s crust (ppm)32027720074016
  • Density : The density of these elements increases down the group as reported below

 

ElementCSiGeSnPb
Density (g/ml)3.51 (for diamond)2.345.327.2611.34
 2.22 (for graphite)    

(4)  Melting point and boiling points

  • The m.pt and b.pt. of this group members decrease down the

 

ElementCSiGeSnPb
m.pt(K)437316931218505600
b.pt.(K)3550312328962024
  • The m.pt and b.pt of group 14 elements are however, higher than their corresponding group 13 elements. This is due to the formation of four covalent bonds on account of four electrons in their valence shells which results in strong binding forces in between their atoms in solid as well as in liquid

(5)  Atomic radii and atomic volume

  • Both atomic radii and atomic volume increases gradually on moving down the group due to the effect of extra shell being added from member to
 CSiGeSnPb
Atomic radius (pm)0.77111122141144
Atomic volume (ml)3.411.413.616.318.27
  • The atomic radii of group 14 elements are than their corresponding group 13 elements due to increase in nuclear charge in the same
  • Some of the ionic radii involving six co-ordination of these group elements are given below,

 

 

 

 CSiGeSnPb
Ionic radius (M2+) in pm73118119
Ionic radius (M++) in pm40536978
  • Electronegativity : The electronegativity decreases from C to Si and then becomes

C        Si       Ge      Sn      Pb Electronegativity on pauling scale 2.5         1.8     1.8     1.7     1.6

The electronegativity from silicon onwards is almost is almost constant or shows a comparatively smaller decreases due to screening effects of d10 electrons in elements from Ge onwards.

(7)  Ionisation energy

  • The ionisation energy decreases regularly down the group; Pb however shows a higher value than Sn due to poor shielding of inner f-orbitals as a result of which effective nuclear charge experienced by outer shell electrons becomes more in
Ionisation energy (kJ mol-1)CSiGeSnPb
IE11086786761708715
IE223521577153714111450
IE346203284330029423081
IE462204354440939294082
  • The first ionisation energies of group 14 elements are higher than their corresponding group 13 elements because of smaller
  • The electropositive character of these elements increases down the group because of decreases in ionisation

(8)  Oxidation state

  • Presence of four electrons in outermost shell of these elements reveals that the members of this family can gain four electrons forming M4+ or M4- ions to show ionic nature or exhibit tetravalent covalent nature by sharing of four electron pairs in order to attain stable
  • The formation of M4+ or M4- ions require huge amount of energy which is normally not available during normal course of reactions, therefore, these elements usually do not form M4+ or M4- ions, but they usually form compounds with covalence of
  • Ge, Sn and Pb also exhibit +2+ oxidation state due to inert pair effect.
  • Sn2+ and Pb2+ show ionic
  • The tendency to form +2 ionic state increases on moving down the group due to inert pair

(9)  Catenation

  • The tendency of formation of long open or closed atom chains by the combination of same atoms in themselves is known as
  • The catenation is maximum in carbon and decreases down the
  • This is due to high bond energy of

 

 

 

  • Only carbon atoms also form double or triple bonds involving pp-pp multiple bond within

> C = C<; – C º C –

  • Carbon also possesses the tendency to form closed chain compounds with O,S and N atoms as well as forming pp-pp multiple bonds with other elements particularly nitrogen and oxygen e.g. C =O; C=N; CºN; C=S are the functional groups present in numerous molecules due to this
  • Carbon can form chain containing any number of carbon atoms Si and Ge cannot extend the chain beyond 6 atoms, while Sn and Pb do not form chains containing more than one or two
  • The reason for greater tendency of carbon for catenation than other elements in the group may further be explained by the fact that the C-C bond energy is approximately of the same magnitude as the energies of the bond between C and other On the other hand, the Si-Si bond is weaker than the bond between silicon and other elements.

 

BondBond energy (k J/mol)BondBond energy (kJ/mol)
C-C348Si-Si180
C-O315Si-O372
C-H414Si-H339
C-Cl326Si-Cl360
C-F439Si-F536

(10)  Allotropy

  • The phenomenon of existence of a chemical element in two or more forms differing in physical properties but having almost same chemical nature is known as If an element or compound exists in two or more forms, it is also known as polymorphism e.g. zinc blende and wurtzite are polymorphs of ZnS. This phenomenon is due to the difference either in the number of atoms in the molecules [as in the case of oxygen (O2) and ozone (O3)] or arrangement of atoms in the molecules in crystal structure (as in the case of various forms of carbon).
  • All the elements of group 14 except lead exhibit
  • Crystalline carbon occurs mainly into two allotropic forms (i) graphite and (ii) diamond (a third allotropic form called fullerenes e.g. C60 and C70 were recently discovered by Prof. Richard E. Smalley and his coworkes), amorphous carbon exists in different forms viz coal, coke, carbon black, lamp black, bone charcoal. Amorphous carbon is usually considered to contain microcrystals of

 

 

 

Carbon

Crystalline                                                              Amorphous (microcrystalline)

 

Diamond            Graphite                                          Coal               Charcoal      Lamp black

 

  • Diamond and Graphite : The two allotropic forms of crystalline Diamond is the purest and hardest form of carbon. Its structure involves a giant molecular form where each carbon atom is surrounded by four other carbon atoms (sp3 hybridization) In doing so, each carbon atom is located in the centre of a regular tetrahedron with its four valencies directed towards the four corners which are linked with four other carbon atoms ( CCC angle = 1090 28’C-C=154 pm = 1.54 Å). The hardness of diamond result due to the uniformity of the C- C covalent bonds. Since the C-C bond length is very small, it has very high density (3.51 g cm-3) and has more compact structure than graphite (density, 2.25 g cm-3) It does not melt (vapourises at 3773K) has very high refractive index (2.45) and is insoluble in all ordinary solvents. It does not conduct electricity as all the four valence electrons are used up in forming covalent bonds with other carbon atoms Diamond, because of its hardness is used in cutting, grinding instruments such as glass and drilling equipments Its ability to reflect and refract light makes diamond an important jewellery material.

Difference between diamond and graphite

 

 

 

Crystalline, transparent with extra brilliance. Hardest form

Bad conductor  of  electricity High Density (3.51 g /cm3) heavy Colourless

Tetrahedral shaped sp3 hybridisation

Less stable, more energy CD ® CG ; DH= – 0.5 k.cal

Used in cutting glass and jewellery; an abrasive

Crystalline, opaque and shiny substance Soft having soapy touch

Good conductor of electricity

Low Density (2.25 g/cm3), lighter than diamond Greyish white

Two dhnensional layer structure having regular hexagonal sheets.

sp2 hybridization

More stable, less energy

CG®CD at high temperature and high P

Used as lubricating agent, electrodes, in pencils, crucibles (due to high m.pt)

 

 

Carbon also exists in three common microcrystalline or amorphous forms (charcoal, carbon black and cocke) Carbon black is formed when hydrocarbons, petroleum, turpentine oil or substances rich in carbon contents are heated in limited supply of oxygen, CH4(g) +O2 (g) ®C(s) + 2H2O(g)

These substances yield a large amount of smoke which is passed into chambers having wet blankets. The soot collected on these blankets is lamp black or carbon black or soot. It is almost pure carbon having as high as 98%to 99% carbon content with small amount of impurities It is a soft black power and is used as a pigment in black inks; large amounts are also used in making automobile tyres.

 

 

 

Charcoal is formed when wood cellulose or other substances containing carbonaceous matter are heated strongly in the absence of air Charcoal has highly open structure, giving it an enormous surface area per unit mass. Charcoal is of various forms such as wood charcoal, sugar charcoal, coconut charcoal, animal charcoal etc. These forms contain varying amount of carbon content. A very pure form of carbon is obtained from sugar. Activated charcoal, a pulverised form whose surface is cleaned. by heating with steam. is widely used to adsorb molecules. It is used in filters to remove offensive odours from air and coloured, foul smelling, bad tasting and toxic chemical as impurities from water.

Coke is an impure form of carbon and is produced when coal is heated strongly in the absence of air (as residue in the destructive distillation of coal) It is widely used as a reducing agent in metallurgical operations.

  • Silicon also exists in crystalline and amorphous allotropic forms Germanium exists in two crystalline allotropic forms Tin has three allotropic forms as grey tin , white tin and rhombic

Graphite occurs in Nature and can also obtained from coke, In graphite, out of four valence electrons, only three form covalent bonds (sp2 hybridization) with three other carbon atoms. This forms hexagonal rings as sheets of on atom thickness. These sheets are held together by weak attractive forces One electron of each carbon atom is free and this enables these thin sheets slide over one another. For this reason graphite is a soft material with lubricating properties.

Graphite is a dark, opaque and soft material (density = 2250 kg/m3) Although graphite is non-metallic still it possesses a metallic lustre. It is insoluble in ordinary solvents. Graphite is a good conductor of heat and electricity because of the present of one free electron on each carbon atom. Graphite is used as a dry lubricant in making electrodes in electric furnaces. It is chiefly used in lead pencils.

Chemical properties

  • Hydrides : All the elements of group 14 combine with hydrogen directly or indirectly to form the covalent

 

hydrides, MH 4

(M = C, Si, Ge, Sn or Pb). The number of hydrides and the ease of preparation decrease on going

 

from carbon to lead.

The hydrides of silicon are called silanes having the general formula Sin H 2n +2 . The hydrides of germanium are called germanes while those of tin are called the stannanes. Only lead forms an unstable hydride of the formula, PbH 4 called the plumbane.

 

Three hydrides of germanium, i.e., GeH 4 ,Ge 2 H 6

and GeH 8

and only two hydrides of tin i.e., SnH 4

and

 

Sn 2 H 6

are well known.

 

  • Oxides : Carbon forms five oxides CO,CO2 ,C 3O2

(carbon suboxide), C5O2

and

C12O9 ,C 3O2

is the

 

anhydride of malonic acid and

CO2

is the anhydride of

H 2CO3

(carbonic acid)

CO2

is a non-polar linear

 

molecule due to maximum tendency of C to form pp–pp multiple bond with oxygen. Si forms SiO2 . Pb forms a

number of oxides. PbO can be obtained by heating Pb(NO3 )2 ,  2Pb(NO3 )2  ¾¾He¾at ® 2PbO + 4 NO2  + O2 .  The

 

red form of PbO is called litharge and the yellow form is massicot.

Pb3O 4 (Red lead, or Sindur) is prepared by

 

heating litharge in air at 470°C, 6PbO + O2

¾¾470¾o¾C ® 2Pb O  ,

Pb3O 4

is a mixed oxide of PbO 2

.2PbO.Pb2O3

 

3
4

is called lead sesquioxide.

GeO2 , SnO2

etc. are also network solids.

 

Note : ® SiO2 ,GeO 2 ,SnO 2 and

PbO2

are all solids.

 

 

 

  • CO2 and SiO2is acidic, GeO 2is weakly acidic while SnO 2 and PbO 2 are amphoteric in

 

  • All the elements of group 14 except silicon from monoxides g., CO,GeO,SnO

these monoxides only CO is neutral, while all other monoxides are basic.

and PbO. Out of

 

  • Halides : Elements of group 14 react with halogens directly to form tetrahedral and covalent halides

 

except C where its halide is produced by the action of halogens on hydrocarbons.

PbBr4

and

PbI 4

do not exist

 

because Pb 4+

is a strong oxidant and

Br

and

I are strong reductants. Hence

Pb 4+

ion is difficult to survive in

 

presence of strong reductants Br and I

and is immediately reduced to

Pb 2+ .

 

Anomalous behaviour of Carbon

Carbon is found to differ in many properties from the rest of the members of group 14. This is because of the following : (i) Its smallest size (ii) Its high electronegativity (iii) Its property to catenate (iv) Absence of d-orbitals in it.

Some of the properties in which it differs from other members are,

  • The melting and boiling points of carbon are very high as compared to the rest to the members of the
  • Carbon in its diamond form is one of the hardest substance
  • It has maximum tendency to show
  • Carbon has high tendency to form Pp – Pp multiple bonds with other elements like nitrogen, oxygen, sulphur Other members of the family form Pp – dp bonds and that also to a lesser extent.
  • CO2 is a gas while the dioxides of all other members are
  • Carbon shows a maximum covalency of four while other members of the family may expand their

 

covalency to six e.g., [SiCl6 ]2 ,[PbCl6 ]2

etc.

 

  • Carbon is not affected by alkalies whereas other members react on For example, silicon form

 

silicates, Si + 2NaOH + 1 / 2 O2 ®

Na2 SiO3 + H 2 .

Sodium silicate

Silicon and its compounds

 

Silicon, being a second member of group – 14, has a much larger size and lower electronegativity than hat of carbon. As a result silicon does not form double bond with itself or with oxygen. Thus SiO bonds are much stronger than Si Si and Si H bonds. Silicon has vacant 3d-orbitals in its valence shell due to which it can extend its covalency from four to five and six.

  • Occurrence : Silicon is the second most abundant element ( 7%) in earth’s crust next to oxygen .It does not occur in free state. It occurs mainly in the form of Silica and silicates. Silicates are formed in rocks and clay

 

as silicates of Mg, Al, K or Fe. e.g. Feldspar ;

K2 Al2 O3 .6SiO2 , Kaolinite;

Al2 O3 .2SiO2 .2H 2 O .

 

  • Preparation : Elemental silicon is obtained by reduction of silica with high purity coke in an electric furnace using excess of silica g. SiO2 + 2C ¾¾® Si + 2CO

 

Very high purity silicon required for making semiconductors is obtained by reduction of highly purified form ( SiHCl3 ) with hydrogen followed by purification by zone refining eg.

SiCl 4

 

SiCl4 + 2H 2 ¾¾® Si + 4 HCl ; SiHCl3 + H 2 ¾¾® Si + 3HCl

 

 

 

 

  • Properties : Silicon exists in three isotopes 14 Si 29

(most common),

14 Si30 with air at high temperature

 

SiO2 form, Si + O2 ¾¾® Si O2 .

With steam, Si reacts when heated to redness to liberate hydrogen,  Si + 2 H 2 O   ¾¾Red¾ne¾ss ® Si O2 + 2 H 2 .

 

With halogens, Si reacts at elevated temperature forming temperature.

SiX4 except fluorine which reacts at room

 

Silicon combines with C at 2500K forming Silicon Carbide (SiC) known as carborundum (an extremely hard

substance),  Si + C  ¾¾250¾0¾K ® SiC.

It reacts with metals like Ca, Mg etc in an electric arc furnace to form Silicides ( Ca2 Si, Mg 2 Si etc.)

 

 

Silicon dissolves in hot aqueous alkalies liberating hydrogen, Si + 4NaOH

¾¾He¾at ®  Na4 SiO4  + 2H 2  ­

 

It also dissolves in fused

Na 2 CO3 displacing carbon

Na2SiO3 +C .

 

  • Uses of silicon : It is added to steel as ferrosilicon ( an alloy of Fe and Si) to make it acid It is used in the pure form as a starting material for production of silicon polymers (Silicones).

(5)  Compounds of silicon

Silica or silicon dioxide ( SiO2 )

It occurs in nature in various forms such as sand, quartz and flint .It is also a constituent of various rocks. It is solid at room temperature. It is insoluble in water.

Silica has a three dimensional network structure in which each Si is bonded to four oxygen atoms which are

 

tetrahedrally disposed around silicon atom. Each O atom is shared by two Si atoms. It may be noted that gas, while SiO2 is hard solid with very high melting point.

CO2 is a

 

Si O2

+ 4HF   ¾® Si F4 + 2H 2 O ; Si F4 + 2HF         ¾®

H2SiF6

(Hydro flouro silicic acid)

 

 

HF readily dissolves Silica, therefore HF can not be store in glass bottles which contain Silica.

It is used in large amount to form mortar which is a building material. It is also used in the manufacture of glass and lenses.

Silicates

4

Almost all rocks and their products (Soil, clay and sand) are made up of silicate minerals and Silica. The basic

 

unit of all silicates is tetrahedral zeolites.

SiO -4 ion. Some of the important silicates are quartz, mica, asbestos, felspar and

 

 

 

 

Silica gel

When a mineral acid (Such as HCl) is added to a concentrated solution of a silicate, gelatinous white ppt. of hydrated silica (silicic acid) separate out.

Na 2 SiO3 + 2HCl ¾¾® 2NaCl + SiO2 .xH 2 O

The white ppt. thus obtained is heated to lose water. When the water content is very low, the solid product is called silica gel. It possesses excellent absorptive properties due to its porous nature and is used for absorbing moisture and an adsorbent in chromatography.

Glass

Glass is an amorphous and transparent solid which is obtained by solidification of various silicates and borates of potassium and calcium.

  • Preparation : Ordinary glass is a mixture of sodium and calcium silicates and is produced by fusing together a mixture of sodium carbonate, calcium oxide and silicon dioxide ( Silica) in a furnace at about 1700K

Na2 CO3 + SiO2 ¾¾® Na 2 SiO3  + CO2 ­ ;  CaO + SiO2 ¾¾® CaSiO3

On continuously heating the entire amount of CO2 is driven out and clear viscous fused mass is obtained. It is poured into moulds to get different types of articles, which are allowed to cool gradually.

This   typed  of  glass  is   called   soda  glass  or   soft  glass  which   has  the   approximate  composition,

Na2 SiO3 , CaSiO3 ,4SiO2 .

  • Various varieties of glass : The different varieties of glasses and their special constituents are given below,

 

Type of glassConstituentsSpecial use
Soft glass Na2CO3, CaCO3, SiO2Ordinary glass for window panes, test tubes, bottles, etc.
Hard glass K2CO3, CaCO3, SiO2For combustion tubes and chemical glassware
High refractive index glassLead oxide, K2CO3For making lenses cut glasses
Pyrex glass Na 2 CO3 , Al 2 O3 , B2 O3 or borax, sandFor high quality glass apparatus cooking utensils
Crook’s glass                         K2CO2, PbCO, CeO2 , sandAbsorbs ultra violet rays, for making lenses
  • Coloured glass : Addition of transition metal compounds to glass give coloured glasses . Small amounts of Cr(III), Mn(IV), Co(II) and Fe(III) compounds impart green, violet blue or brown colour respectively

 

Nitrogen is the first member of group 15 or VA of the periodic table. It consists of five elements nitrogen (N), phosphorus (P), arsenic (As), antimony (Sb) and bismuth (Bi). The elements of this group are collectively called

 

 

 

pnicogens and their compounds as pniconides. The name is derived from Greek word “Pniomigs” meaning suffocation. Pniconide contain M 3 species.

Electronic configuration

 

Elements        Electronic configuration (ns 2 np3)
7 N                        1s 2 ,2s 2 2p3 or [He] 2s 2 2p 3

15 P                        1s 2 ,2s 2 2p6 ,3s 2 3p3 or [Ne] 3s 2 3p 3

33 As                     1s 2 ,2s 2 2p6 ,3s 2 3p6 3d10 ,4s 2 4 p3 or [Ar] 3d10 4s 2 4 p 3

51 Sb                      1s 2 ,2s 2 2p6 ,3s 2 3p6 3d10 ,4s 2 4 p6 4d10 ,5s 2 5p 3 or [Kr] 4d10 5s 2 5p 3

83 Bi                      1s 2 ,2s 2 2p6 ,3s 2 3p6 3d10 ,4s 2 4 p6 4d10 4 f 14 ,5s 2 5p6 5d10 ,6s 2 6p 3 or [Xe] 4 f 14 5d10 6s 2 6p 3

Physical properties

  • Physical state : Nitrogen– (gas), phosphorus – (solid) (vaporises easily), As, Sb, Bi–solids.

Note : ® Nitrogen is the most abundant gas in the atmosphere. It constitutes about 78% by volume of the atmosphere. Phosphorus is the most reactive element in this group and its yellow form is always kept under water.

  • Atomic radii : Atomic radii increases with atomic number down the group e., from N to Bi due to addition of extra principal shell in each succeding elements.
  • Ionisation energy : The ionisation values of the elements of this group decreases down the group due to gradual increases in atomic
  • Electronegativity : Generally the elements of nitrogen family have high value of This value shows a decreasing trend in moving down the group from nitrogen to bismuth.
  • Non-metallic and metallic character : Nitrogen and phosphorus are non-metals, arsenic and antimony are metalloids (semi-metal) and bismuth a typical
  • Molecular state : Nitrogen readily forms triple bond (two pp –pp bonds) and exists as discrete diatomic

 

gaseous molecule

(N º N)

at room temperature. Phosphorus, arsenic and antimony exist in the form of discrete

 

tetra atomic molecules such as P4 , As4 , Sb4

in which the atoms are linked to each other by single bonds.

 

  • Melting and boiling points : The melting points and boiling points of group 15 elements do not show a regular

Note : ® M.pt. first increases from N to As and then decreases from As to Bi. Boiling point first increases from N to Sb. Boiling point of Bi is less than Sb.

  • Allotropy : All the members of group 15 except Bi exhibit the phenomenon of
    • Nitrogen exists in two solid and one gaseous allotropic forms.
    • Phosphorus exists in several allotropic forms such as white, red, scarlet, violet and black

 

  • White or yellow phosphorus : White phosphorus is prepared from rock phosphate coke which are electrically heated in a

Ca3 (PO4 )2 , SiO2

and

 

2Ca3 (PO4 )2  + 6SiO2  ¾¾D ® 6CaSiO3  + P4 O10 ;

When exposed to light, it acquires a yellow colour.

P4 O10  + 10C ¾¾D ® P4  + 10CO

 

 

 

  • Red phosphorus : It is obtained by heating yellow phosphorus, between 240 –250°C in the presence of an inert gas. Yellow phosphorus can be separated from red phosphorus by reaction with NaOH (aq) or KOH (aq) when the former reacts and the latter remains
    • Arsenic exists in three allotropic forms namely grey, yellow and black. Antimony also exists in three forms, , metallic, yellow and explosive.
  • Oxidation state : The members of the group 15 exhibit a number of positive and negative oxidation
  • Positive oxidation states : The electronic configuration (ns 2np3 ) for the valence shell of these elements shows that these elements can have +3 and +5 oxidation In moving down this group, the stability of +3 oxidation state increases. It may be pointed out here that nitrogen does not exhibit an oxidation state of +5,

because it fails to expand its octet due to nonavailability of vacant d-orbitals.

  • Negative oxidation states : For example oxidation state of nitrogen is –3. The tendency of the elements to show –3 oxidation state decreases on moving down the group from N to Bi.
  • Catenation (self linkage) : Elements of group 15 also show some tendency to exhibit This tendency goes on decreasing in moving down the group due to gradual decrease in their bond (MM) energies.

Note : ® Out of the various allotropic forms of phosphorus, black phosphorus is a good conductor of electricity (similarity with graphite).

  • Proteins, the building blocks of our body contain 16% of nitrogen in
  • Radioactive phosphorus (P 32 ) is used in the treatment of leukemia (blood cancer).
  • The disease caused by the constant touch with white phosphorus is called Phossy

Chemical properties

 

  • Hydrides : All the members form volatile hydrides of the type

AH3 . All hydrides are pyramidal in shape.

 

The bond angle decreases on moving down the group due to decrease in bond pair–bond pair repulsion.

 

NH3 107o

PH 3

94o

AsH 3

92o

SbH3

91o

BiH3

90o

 

The decreasing order of basic strength of hydrides is as follows :

NH3  > PH 3  > AsH 3  > SbH3  > BiH3 .

 

The increasing order of boiling points is as follows : PH 3 < AsH 3 < NH 3 < SbH 3 .

 

NH 3

is thermally most stable and

BiH3

is least stable. This is because in

NH3 , N H covalent bond is the

 

strongest due to small size of N atom. Hence, the decomposition temperature of

NH 3

will be the highest.   The

 

increasing order of reducing character is as follows,

NH 3 < PH 3 < AsH 3 < SbH 3 < BiH 3 .

 

Note : ® Diphosphine (P2 H4 ) and hydrazine (N 2 H4 ) are other two important hydrides. Hydrazine a strong reducing agent, is used in organic synthesis and rocket fuels and is prepared as follows, 2NH 3 + NaOCl ® N 2 H4 + NaCl + H 2 O

  • Phosphine is poisonous and does not form any

 

  • Phosphine forms vortex rings of

PO5

in the form of white smoke when it comes in contact with air

 

due to combustion. This is due to impurities of diphosphine (P2 H4 ) .

  • Phosphine is used to prepare smoke screens in Calcium phosphide reacts with water to

 

form phosphine which burns to give clouds of

  • Liquor ammonia is a concentrated solution of

PO5

NH 3

which acts as smoke screens. in water.

 

  • Ammonia will not burn in air, but burns in pure O 2 with a yellowish flame to produce N 2 and H2O .

 

 

 

 

  • Halides : The members of the family form trihalides (MX 3 )

and pentahalids

(MX5 ) . The trihalides are

 

sp3 -hybridized with distorted tetrahedral geometry and pyramidal shape while pentahalides are sp3 d-hybridized

and are trigonal bipyramidal in shape. The trihalides are hydrolysed by water and ease of hydrolysis decreases

 

when we move down the group. Hence,

NCl3

is easily hydrolysed but

SbCl3

and

BiCl3

are partly and reversibly

 

hydrolysed.

NF3

is not hydrolysed due to lack of vacant d-orbital with nitrogen.

PF3

and

PF5

are also not

 

hydrolyzed because the P – F bond is stronger than P – O covalent bond. The hydrolysis products of the other

 

halides are as follows :

NCl3  + 3H 2 O ® NH 3  + 3HOCl  ;

PCl3  + 3H 2 O ® HPO3  + 3HCl

 

2AsCl3 + 3H2O ® As2O3 + 6HCl   ;

SbCl3  + H 2 O ® SbOCl + 2HCl  ;

BiCl3 + H 2 O ® BiOCl + 2HCl

 

Their basic character follows this decreasing order as

NI 3 > NBr3 > NCl3 > NF3 . Except

NF3 , the trihalides

 

of nitrogen are unstable and decompose with explosive violence.

NF3

is stable and inert.

NCl3

is highly explosive.

 

Trifluorides and trichlorides of phosphorus and antimony act as Lewis acid. The acid strength decreases down the

 

group. For example, acid strength of tri-chlorides is in the order ;

PCl3 > AsCl3 > SbCl3 .

 

Nitrogen does not form pentahalides due to non-availability of vacant d-orbitals. The pentachloride of phosphorus is not very stable because axial bonds are longer (and hence weaker) than equitorial bond. Hence,

 

PCl5

decomposes to give

PCl3

and Cl2 ;

PCl5  ⇌

PCl3  + Cl2 .

 

The unstability of

PCl5

makes it a very good chlorinating agent. All pentahalides act as lewis acids since they

 

can accept a lone pair of electron from halide ion.

 

Note : ® Solid

PCl5

is an ionic compound consisting of

[PCl4 ]+

[PCl6 ] ,

[PCl4 ]+

has a tetrahedral

 

structure, while [PCl6 ]

has an octahedral structure.

 

  • Since,

PCl5

reacts readily with moisture it is kept in well stoppered bottles.

 

  • PI 5

does not exist due to large size of I atoms and lesser electronegativity difference between

 

phosphorus and iodine.

  • Down the group, the tendency to form pentahalides decreases due to inert pair e.g.,

does not exist.

BiF5

 

  • Oxides : These elements form oxides of the type

XO3 , XO4

and

X2O5  .

 

  • Oxides of Nitrogen : Nitrogen forms two more oxides e., N 2O

and NO and both are neutral. Nitrous

 

oxide ( N 2O ) has a sweet taste and its main use is as anaesthetic. When inhaled in mild quantities it causes hysterical laughter so it is also called Laughing gas. Nitric oxide (NO) can be obtained by treating a mixture of

 

sodium nitrite and ferrous sulphate with dil.

H 2 SO4 . Other oxides of nitrogen are :

NO2 ,

N 2 O3 ,

N 2 O5 .

 

The acidic strength of oxides :

N 2 O < NO < N 2 O3  < N 2 O4  < N 2 O5 .

 

  • Oxides of phosphorus :

PO6

(Phosphorus trioxide),

P4 O10 (Phosphorus pentaoxide).

 

  • Oxides of other elements : The decreasing order of stability of oxides of group 15 follows as,

P2 O5 > As2 O5  > Sb2 O5  > Bi2 O5

 

Except

P2O5 , all pentaoxides show oxidising properties. Also

PO5

is acidic in nature.

NO5

is the strongest

 

oxidising agent. The nature of oxides of group 15 elements is as follows,

 

N 2 O3

and

PO3

(acidic) ;

As2O3

and Sb2 O3 (amphoteric) ;

Bi2 O3

(basic)

 

  • Oxyacids : Oxyacids of nitrogen are

HNO2 , HNO3 , HN 2 O4

(Nitroxylic acid)

and

HNO4

(Pernitric acid)

, which are explosive.

 

 

 

 

Note   : ®

HNO3

is called aqua fortis and prepared from air (Birkel and Eyde process) and

NH3

 

(Ostwald process). It acts as a strong oxidising agent.

Oxyacids of phosphorus are,

  • H3 PO2 (Hypophosphorus acid) : Reducing agent and

 

HPO3

HPO4

(Orthophosphorus acid) : Reducing agent and dibasic. (Orthophosphoric acid) : Weak tribasic acid.

 

HPP7

(Pyrophosphoric acid) : It is obtained by heating

HPO4

to 220°C. It is tetrabasic.

 

HPO3

(Metaphosphoric acid) : It is formed by the dehydration of

HPO4

at 316°C. Also exists as a trimer

 

and is monobasic.

  • H4PO6 (Hypophosphoric acid) : Tetrabasic
  • H4 P2 O5 (Pyrophosphoric acid) : Dibasic acid

Anamalous behaviour of Nitrogen

Nitrogen is known to differ form other members of the family because of the following facts,

  • Its small size (ii) Its high electronegativity (iii) Its high ionisation energy (iv) non-availability of d-orbital in the valence (v) Its capacity to form pppp multiple bonds.

The main points of difference are,

  • Nitrogen is a gas (N 2 ) while other members are
  • Nitrogen is diatomic while other elements like phosphorus and arsenic form tetra-atomic molecules (P4 , As4 ) .

 

  • Nitrogen form five oxides two oxides (tri and pentaoxides).

(N 2 O, NO, N 2 O3 , N 2 O4

and

NO5 )

while other members of the family form

 

  • Hydrides of nitrogen show H-bonding while those of other elements do
  • Nitrogen does not show pentacovalency because of absence of d-orbitals while all other elements show
  • Nitrogen dos not form complexes because of absence of d-orbitals while other elements show complex

 

formation e.g., [PCl6 ] ,[AsCl6 ]

etc.

 

  • The hydride of nitrogen (NH 3 ) is highly basic in nature while the hydrides of other elements are slightly

 

  • Except for

NF3 , other halides of nitrogen e.g.,

NCl 3 , NBr3

and

NI 3

are unstable while the halides of

 

other elements are fairly stable.

Nitrogen and its compounds

 

N 2 was discovered by Daniel Rutherford. It is the first member of group 15 in the periodic table.

 

(1)   Occurrence :

N 2 , occurs both in the free state as well as in the combined state.

N 2 occurs in

 

atmosphere to the extent of 78% by volume in free state.

N 2 is present in many compounds such as potassium

 

nitrate (nitre). Sodium nitrate (Chile salt peter) and many ammonium compounds. of proteins in plants and animals in combined state.

  • Preparation : It is prepared by the following methods,

N 2 is an important constituent

 

 

 

 

  • Laboratory method : In the laboratory

N 2 is prepared by heating an aqueous solution containing an

 

equivalent amounts of

NHCl

and

NaNO2 .

 

NH 4 Cl (aq.) + NaNO2 (aq.) ¾¾He¾at ® N 2 (g) + 2H 2 O(l) + NaCl

  • Commercial preparation : Commercially N 2 is prepared by the fractional distillation of liquid

 

(3)  Physical properties :

N 2 is a colourless, odourless and tasteless gas. It is a non-toxic gas. It’s vapour

 

denstiy is 14. It has very low solubility in water.

(4)  Chemical properties

  • N 2 is neutral towards It is chemically unreactive at ordinary temp. It is neither combustible nor it

 

supports combustion.

  • The N – N bond in

N 2 molecule is a triple bond

(N º N)

with a bond distance of 109.8 pm and bond

 

dissociation energy of 946 kJ mol-1

  • Combination with compounds :

N 2 combines with certain compounds on strong heating . eg

 

CaC2

+ N 2  ¾¾130¾0¾K ® CaCN 2  + C  ;

Al 2 O3 + N 2  + 3C ¾¾210¾0¾K ® 2AlN + 3 CO

 

Calsium carbide

Calsium cyanamide

Aluminium oxide

Al. nitride

 

Both these compounds are hydrolysed on boiling with water to give ammonia.

CaCN 2 + 3HO ¾¾® CaCO3 + 2NH3 ; AlN + 3H 2 O ¾¾® Al (OH)3 + NH 3

Therefore, calcium cyanamide is used as a fetilizer under the name nitrolim (CaCN 2 + C )

 

(5)  Uses of nitrogen :

N 2 is mainly used in the manufacture of compounds like

NH 3 , HNO3 , CaCN 2

etc.

 

(6)  Compounds of nitrogen

  • Hydrides of nitrogen

 

Ammonia

 

Ammonia is the most important compound of nitrogen. It can be manufactured by Haber’s process. In this

 

process, a mixture of

N 2 and

H 2 in the ratio of 1 : 3 is passed over heated Fe at 650 –800K as catalyst and Mo as

 

promotor,

N 2 + 3H 2 ⇌

2NH 3 , DH = -93.6 kJ mol -1

This is a reversible exothermic reaction.

 

Ammonia is prepared in the laboratory by heating ammonium salt (NH 4 Cl) with a strong alkali like NaOH

 

NHCl + NaOH ¾¾® NH 3 + HO + NaCl

Ammonia can be dried by passing over quick lime (CaO). How

 

ever, it can not be dried with dehydrating agents such as conc. ammonia reacts with these compounds.

H 2 SO4 , P2 O5

and anhydrous

CaCl 2

because

 

NH 3

is a colourless gas with a characteristic pungent smell called ammonical smell. It is highly soluble in

 

water and its solution is basic in nature,

NH 3 + H 2 O

NH + + OH

 

4

NH 3

is expected to have a tetrahedral geometry, but the lone pair distorts its geometry and the molecule has

 

pyramidal geometry with N – H bond length of 101.7 pm and a bond angle of 107.5o. Liquid ammonia is widely used as a refrigerant due to its high heat of vaporization.

Hydrazine, (NH2 – NH2)

Hydrazine is prepared commercially by boiling aqueous ammonia or urea with sodium hypochloride in the presence of glue or gelatin.

2NH3 + NaOCl  ® N2H4 + NaCl + H2O

 

 

 

The resulting solution is concentrated and anhydrous hydrazine may be obtained by further distillation over barium oxide. Alternatively, the hydrazine present in the resulting solution is precipitated as sparingly soluble crystalline hydrazine sulphate on treatment with sulphuric acid, NH2NH2 + H2SO4 ® N2H4.H2SO4.

The precipitate is removed and treated with an alkali when hydrazine hydroxide H2N.NH3OH is obtained.

This is distilled under reduced pressure, over barium oxide to liberate free hydrazine.

H2N.NH3OH    + BaO     ¾¾dis¾til ®  NH2NH2   + Ba(OH)2

Physical properties : Anhydrous hydrazine is a colourless fuming liquid (m.p. 20C and b.p. 1140C ) soluble in water in all proportions. It is also soluble in alcohol. It is strongly hygroscopic.

Chemical properties : It behaves as a diacid base, Thus with hydrochlorides it forms hydrazine monochloride H2N.NH3Cl and hydrazine dichloride ClH3N.NH3Cl.

 

  • Hydrazine burns in air with the evolution of N2H4 + O2 ® N2 + 2H2O The alkyl derivative of hydrazine are used as rocket fuels.
  • It reduces Fehling’s solution to red cuprous oxide and iodates to

DH = -622 kJ

 

3

4Cu2+  +  NH2NH4    ® 4Cu+ + 4H+ + N2 ; 2lO + 3NH2NH2  ® 2l +  6H2O +  3N2

Uses : Hydrazine is used as a rocket fuel. It is also used as a reagent in organic chemistry.

 

  • Oxides of nitrogen : Nitrogen combines with

O2 under different conditions to form a number of binary

 

oxides which differ with respect to the oxidation state of the nitrogen atom. The important oxides are

 

N 2 O, NO, N 2 O3 , NO2 , N 2 O4

and

N 2 O5 .

Oxides of Nitrogen

 

Oxide                    Oxidation State of N               Physical appearance                               Structure
Nitrous oxide ( N 2O )+1Colourless gasN º N ® O

N = O

O    N N

O

N N

O

 

N

O

O         O

N     N

O

  
Nitric oxide (NO)+2Colourless  
Dinitrogen trioxide ( N 2 O3 )+3Blue solidO 
   O 
Dinitrogen                tetraoxide+4Colourless liquidO 
( N 2 O4 )  O 
Nitrogen dioxide ( NO2 )+4Brown gas  
   O 
Dinitrogen                pentoxide ( N2O5 )+5Colourless gas O

O

 

 

 

 

 

  • Oxyacids of nitrogen

Oxyacids of nitrogen

 

Name of oxoacidM. F.StructureOxidation State of NBasicitypKaNature
Hyponitrous acidH2N2O2 

H – N = O

¯

O

H O N = O

¯

O

O = N- O – O – H

¯

O

+12(dibasic)Very weakHighly explosive
Nitrous acidHNO2+31

(monobasic)

3.3Unstable Weak acid
Nitric acidHNO3+51

(monobasic)

-3.0Stable, Strong acid
Pernitric acidHNO4+51

(monobasic)

 Unstable and explosive

Phosphorus and its compounds

It is the second member of group 15 (VA) of the Periodic table. Due to larger size of P, it can not form stable P p – Pp bonds with other phosphorous atoms where as nitrogen can form Pp – Pp bonds .

  • Occurence : Phosphorous occurs mainly in the form of phosphate minerals in the crust of Some

 

of these are : (i) Phosphorite

Ca3 (PO4 )2 , (ii) Fluorapatite

Ca5 (PO4 )3 F , (iii) Chlorapatite 3 Ca3 (PO4 )2 .CaCl2 ,

 

  • Hydroxyapatite;

Ca5 (PO4 )3 OH . Phosphates are essential constituents of plants and animals . It is mainly

 

present in bones, which contains about 58% calcium phosphate.

  • Isolation : Elemental phosphorus is isolated by heating the phosphorite rock with coke and sand in an

 

electric furnace at about 1770K,

2Ca3 (PO4 )2 + 6SiO2 ¾¾® 6CaSiO3 + P4 O10 ;

Calicum silicate

PO10 + 10C ¾¾® P4 + 10CO

 

  • Allotropic forms of phosphorus : Phosphorus exists in three main allotropic forms,
    • White phosphorus, (ii) Red phosphorus, (iii) Black phosphorus

 

Some physical properties of three forms of phosphorus

PropertiesWhite phosphorusRed phosphorusBlack phosphorus
ColourWhite but turns yellow on exposureDark redBlack

 

 

 

 

 

Compounds of phosphorus

(1) Oxides and oxyacids of phosphorus : Phosphorus is quite reactive and forms number of compounds in oxidation states of –3 , +3 and +5.

  • Oxides : Phosphorus forms two common oxides namely, (a) phosphorus trioxide ( P4 O6 ) and (b) phosphorus penta oxide (P4 O10 ) .

 

  • Phosphorus (III) oxide ( P4 O6 ) : It is formed when P is burnt in a limited supply of air, P4 + 3O2

(limited)

It is a crystalline solid with garlic odour. It dissolves in cold water to give phosphorous acid,

® P4 O6 .

 

P4O6 + 6H 2 O

cold

®   4 H3 PO3

Phosphorous acid

, It is therefore, considered as anhydride of phosphorus acid.

 

With hot water, it gives phosphoric acid and inflammable phosphine, P4O6 + 6H2O (hot) ®

3H 3 PO4  + PH 3

Phosphoric acid

 

It reacts vigorously with Cl2 to form a mixture of phosphoryl chloride and meta phosphoryl chloride.

 

P4 O6 + 4Cl2  ®

2POCl3     +

Phosphoryl chloride

2PO2 Cl

Metaphosphoryl chloride

 

  • Phosphorus  (V)   oxide    (P4O10 ):  It  is  prepared  by  heating  white  phosphorus  in  excess  of  air,

P4 + 5O2 (excess) ® P4 O10 . It is snowy white solid. It readily dissolves in cold water forming metaphosphoric acid.

 

P4 O10 + 2H 2 O ®

4 HPO3

. With hot water, it gives phosphoric acid,

P4 O10  + 6H 2 O ® 4 H 3 PO4   .

 

(Cold)

Metaphosphoric acid

Het

Phosphoric acid

 

P4O10

is a very strong dehydrating agent. It extracts water from many compounds including

HSO4 and

HNO3 ,

 

H 2 SO4  ¾¾P4 O¾10  ® SO3 ; 2HNO3  ¾¾P4 O¾10  ® N2O5 ; CH3CONH 2 ¾¾P4 O¾10  ®  CH3CN

 

H2O

H 2O

Acetamide

H2O

Methyl cyanide

 

  • Oxyacids of phosphorus : Phosphorus forms a number of oxyacids which differs in their structure and oxidation state of

Oxyacids of phosphorus

 

 

 

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